GIFT   ©F 
Harry  East  Miller 


A  SHORT  COURSE 


IN 


QUALITATIVE  ANALYSIS 


BY 

J.  M.  CRAFTS, 

FORMERLY  PROFESSOR   OF   GENERAL   CHEMISTRY  IN   THE   CORNELL    UNIVERSITY. 


THIRD  EDITION. 


REVISED   BY 


CHARLES  A.  SCHAEFFER,  A.M.,  PH.D., 

PROFESSOR  OF   GENERAL   AND   ANALYTICAL  CHEMISTRY  IN 
THE  CORNELL  UNIVERSITY. 


NEW  YORK: 

JOHN    WILEY    &    SONS, 

15  ASTOR  PLACE. 

1880. 


COPYRIGHT, 

1879, 
BY  JOHN  WILEY  &  SONS. 

GIFT  OF 


i-ffo 


This  Book  is  inscribed  by  the  Author 


M.  ADOLPH  WURTZ, 


AS  A  TOKEN  OF  AFFECTIONATE  REGARD,   TO    A    FRIEND    AND    TEACHER, 

AND  AS  A  TRIBUTE  OF  RESPECT  AND  ADMIRATION,   TO  THE 

MASTER  OF  A  SCHOOL  IN  MODERN  CHEMISTRY. 


M81896 


PREFACE  TO  THE  THIRD  EDITION. 


IN  preparing  a  new  edition  of  this  work,  the  editor  has  made 
no  essential  alteration  of  the  plan  adopted  by  Professor  Crafts 
in  the  previous  editions.  He  has  endeavored,  however,  as  far 
as  possible,  to  embody  the  practical  results  of  the  great  ad- 
vance made  in  both  theoretical  and  practical  chemistry  within 
the  past  ten  years.  The  few  important  changes  thereby  ren- 
dered necessary  may  be  briefly  specified. 

Chapter  II.,  on  chemical  nomenclature,  has  been  entirely 
rewritten.  The  usefulness  of  the  work  has  been  somewhat  ex- 
tended by  the  introduction  of  several  elements  not  included 
in  the  previous  editions,  viz.  :  strontium,  cadmium,  iodine,  and 
bromine. 

Further  changes  have  been  made  in  the  methods  of  separa- 
tion and  detection  of  several  of  the  elements ;  more  reliable 
processes  being  substituted  for  those  which  in  the  hands  of 
students  were  found  not  to  be  entirely  satisfactory. 

The  editor  desires  especially  to  acknowledge  his  obligation 
to  his  friend  and  colleague,  Dr.  G.  C.  Caldwell,  for  numerous 
valuable  suggestions  of  which  he  has  availed  himself. 

ITHACA,  December  17,  1879. 


PREFACE  TO  THE  FIRST  EDITION. 


THIS  little  work  was  written  for  the  use  of  a  class  of  stu- 
dents in  the  Cornell  University,  who  take  a  year's  course  of 
chemistry,  including  four  hours  a  week  of  laboratory  practice  ; 
reference  was  also  had  to  the  requirements  of  the  scientific 
students  in  Union  College,  whose  course  is  nearly  equivalent 
to  that  mentioned.  The  author  is  indebted  to  the  kindness 
of  Professor  Perkins  for  many  valuable  suggestions,  and  for 
the  compilation  of  Tables  IV.  and  V.  at  the  end  of  the  book. 

A  considerable  portion  of  the  introductory  part  of  this  book 
is  devoted  to  an  explanation  of  the  theory  of  chemical  reac- 
tions and  nomenclature.  Many  of  the  standard  works  on  an- 
alytical chemistry  still  use  the  old  notation,  and  the  formulas 
to  be  found  in  them  do  not  correspond  to  those  used  in  the 
best  text-books  on  general  chemistry  ;  and  none  of  the  ele- 
mentary works  for  the  laboratory-student  supply  the  want, 
often  felt  by  him,  of  a  system  of  rules,  at  hand  for  use  at  the 
moment  when  he  most  requires  them,  namely,  when  he  is 
writing  the  formula  of  a  reaction  at  his  desk  in  the  laboratory 
with  his  tests  before  him.  The  author  has  had  these  points 
in  view  in  writing  the  two  introductory  chapters. 

It  might  be  objected  that  the  theory  of  chemical  notation 
should  be  found  in  text-books  on  general  chemistry  ;  but  even 
when  the  student  has  mastered  the  rudiments  of  the  science, 
as  given  in  any  of  the  best  modern  works,  he  will  find  the 


PREFACE.  vii 

arrangement  of  many  of  them  inconvenient  for  reference, 
although  it  is  excellent  for  instruction,  and,  moreover,  it  is  by 
no  means  necessary  that  the  study  of  those  works  should  pre- 
cede that  of  analytical  chemistry. 

It  is  quite  feasible  to  commence  a  course  of  experiments  in 
the  laboratory  concurrently  with  the  study  of  any  of  the  ele- 
mentary text-books  on  general  chemistry,  or  with  the  attend- 
ance upon  lectures,  illustrated  with  the  usual  experiments. 
Chemistry  certainly  becomes  a  more  attractive  study  when 
the  practical  and  the  theoretical  present  themselves  side  by 
side,  so  that  while  the  theory  explains  the  experiment,  the  ex- 
periment awakens  an  interest  in  the  theory  ;  and  no  course  of 
study  is  more  apt  to  interest  the  beginner  in  chemistry  than 
that  of  the  admirably  simple  and  delicate  tests  of  qualitative 
analysis  ;  tests  which  illustrate  the  general  laws  of  the  science, 
while  they  have  a  very  direct  bearing  upon  some  of  the  prob- 
lems of  every-day  life. 

Analytical  chemistry,  besides  its  immediate  value  as  an  im- 
portant branch  of  knowledge,  cannot  be  too  highly  prized  as 
affording  a  convenient  introduction  to  the  methods  of  investi- 
gation used  in  an  experimental  science,  and  as  offering  a 
means  of  education  of  many  faculties,  which  are  not  easily 
developed  by  school  or  university  training.  The  importance 
of  laboratory  experiments  is  awakening  every  day  increased 
attention,  and  the  time  is  fast  passing  by  when  chemistry  is 
taught  to  persons,  who  suppose  that  they  have  a  vocation  for  a 
scientific  profession,  only  by  lectures  and  recitations. 

The  system  of  analysis,  given  in  Part  III.,  is  founded  upon 
that  of  Fresenius,  and  includes  a  minute  description  of  all  the 
steps  to  be  taken  in  performing  tests  and  directions  for  pass- 


viii  PREFACE. 

ing  from  one  test  to  another  ;  and  since  these  details,  which 
are  required  by  the  beginner,  not  only  become  unnecessary, 
after  a  certain  familiarity  with  analytical  work  has  been  ac- 
quired, but  also  render  the  scheme  less  convenient  for  refer- 
ence, tables  similar  to  those  of  Will  have  been  added,  which 
indicate  the  important  tests  and  leave  the  experience  of  the 
student  to  suggest  the  proper  mode  of  applying  them. 

The  Tables  IV.  and  V.  are  intended  to  record  a  number  of 
facts  in  analytical  chemistry  in  a  compact  form,  and  to  give 
an  exact  conception  of  what  is  meant  by  the  insolubility  of  a 
precipitate,  and  a  means  of  judging  of  the  advantages  of  dif- 
ferent methods  of  precipitation. 

ITHACA,  August  19,  1869. 


PREFACE  TO  THE  SECOND  EDITION. 


A  FEW  changes  have  been  made  in  the  second  edition, 
which  have  been  suggested  by  the  use  of  the  previous  one  in 
the  laboratory  and  some  practical  directions  have  been  added, 
which  are  intended  to  mark  out  more  closely  the  course  of 
study  which  it  seems  most  advisable  to  follow. 

One  of  the  principal  applications  of  analytical  chemistry  is 
to  the  investigation  of  minerals,  and  it  may  not  be  out  of  place 
here  to  recommend  to  the  consideration  of  instructors  and 
students  the  combination  of  mineralogical  work  with  labora- 
tory practice.  A  convenient  order  of  study  is,  first  to  acquire 


PREFACE.  ix 

familiarity  with  the  methods  of  detection  and  separation  of 
bodies  in  solution,  which  are  given  in  Part  III.,  and  then,  be- 
fore taking  up  the  analysis  of  solid  bodies,  to  study  their  crys- 
tallographical  and  general  mineralogical  characters  from  any 
elementary  work  treating  of  the  subject.  Turning  again  to 
the  general  analysis  of  solid  bodies  (minerals  and  technical 
products  of  all  kinds),  the  course  should  be  made  as  varied 
and  extended  as  possible  ;  and  it  will  be  found  that  the  drill 
in  blowpipe  practice  and  the  study  of  the  physical  characters 
of  minerals  make  an  excellent  preparation  for  the  preliminary 
testing  of  solids,  Part  III.,  and  often  afford  hints  which  are  of 
great  value  in  the  subsequent  performance  of  the  complete 
analysis  in  the  wet  way. 

It  is  quite  possible  for  a  class  to  complete  a  course  like  the 
one  proposed  by  working  four  hours  a  week  for  a  year  in  the 
laboratory. 

ITHACA,  1870^ 


A  SHORT  COURSE 

IN 

ANALYTICAL  CHEMISTRY. 

• 
PART  L— INTRODUCTION. 

CHAPTER   I 

THE  analytical  chemist  is  not  generally  required  to  investi- 
gate a  large  range  of  substances,  and  the  theory  of  the  com- 
position of  those  bodies  which  he  usually  deals  with  can  be 
briefly  explained. 

Elements.  The  ultimate  result  of  the  analysis  or  de- 
composition of  all  matter  is  the  discovery  of  a  number  of  sub- 
stances which  cannot  be  decomposed,  and  are  therefore  called 
simple  or  elementary  substances.  The  table  on  the  next  page 
contains  the  names  of  thirty-six  elements.  They  are  chosen 
from  among  the  seventy  or  more  elements  which  have  been 
discovered,  because  they  are  more  frequently  met  with  than 
the  others,  and  because  they  are  the  only  ones  treated  of  in 
this  work.* 

*  Within  the  past  four  years  ten  new  elements  have  been  reported  by 
their  discoverers.  How  many  of  these  will  be  substantiated  it  is  at 
present  impossible  to  say.  On  the  other  hand,  very  recent  investiga- 
tions seem  to  show  that  at  least  several  of  our  well-known  elements  are 
really  compound  bodies.  The  exact  number  of  elements  is,  therefore, 
especially  at  the  present  crisis,  very  difficult  to  fix. 
I 


. 

•-                                                                   »n  N 

*     *    *     * 

*  ^-^         *       fc*fc 

*                                                               m  vn 

1'"' 

*:V':  •'"'"' 

Id 

X 

C/2 

§  g" 

w 

^ 

pj  _2 

E 

1 

So  S 

,£" 

S  8 

T$-  M 
H    CO 

in  ci  co* 

Cfl 

M    M 

Q 

'PH' 

*  ^*      * 

£j 

^*  t/T 

4j  w^S 

W 

g| 

o  §-5 

PH 

g5^ 

ill 

gl 

<<S 

co  eJ 

Tt-                                         •* 

N    M 

C>             M                                    »n  in  m  W 

M                    M 

Cfl 

•       * 

Q 

g     J^* 

P< 

PH                                                  <u    -  a 

H 

c^  ^ 

~                                                                             0)    ^    § 

W 

§1 

§    «s           ^111 

.a-2 

'"S               C*                                     §   rt         § 

rr5  ^ 

S                   H                                                   HH  1^  U  < 

rf  M 
M    H 

xn  0*  co        O 

^28     2* 

1 

!?  : 

"^s*   i 

2. 

fTW 

^**  '  •        ^ 

o"  o  "5     ^ 

H 

f! 

II 

^  ?  .52      ? 
<j  <^  pq      ^5 

^-                                               <o 

O    « 
M    CO 

O  O   O  M         in  *n  in  mvO         W  ^*co  co 

t 

C4          C4    M                                                                     M 

. 

• 

C/3 

Q 

J 

:  tDQ      :  :    ^       ^  j  ^  : 

£ 

(Dc« 

xi'SfliscS     «       §u  g>(^ 

Q 

i 

O  C/3 

||1|    llg^g     Illl 

<U   O  ulj    d        rj*^    O    ti  w—  i  »F^        UHI    w  -Ij    ™ 

H-3U^U      £(Jh-5SN      SUc^PQ 

in 

W    COCO    N 

CO                                                                                       CO  O^ 

C/2 

M 

M 

Q 

*?A  i  * 

•i                                        Irf 

O 

.3.2  .5  jg1 

M     M     0     A     b 

Islfc 

I                                       U 

SYMBOLS.  3 

The  task  of  the  analytical  chemist  is  to  recognize  these  ele- 
ments when  they  occur  alone,  or  to  separate  them  from  the 
compounds  of  which  they  form  a  part,  or  most  frequently  to 
resolve  a  compound  substance  into  simpler  compounds,  and 
to  isolate  the  latter  in  such  form  that  they  can  be  easily  re- 
cognized. When  the  different  elements  are  known  which 
constitute  such  simpler  compounds,  the  composition  of  the 
complex  body  from  which  they  were  obtained  can  be  deter- 
mined. 

All  compound  substances  are  formed  by  the  union  of  chem- 
ical elements,  and  a  previous  study  of  the  manner  in  which 
the  elements  combine  with  each  other  is  essential  to  the  suc- 
cessful pursuance  of  analytical  investigations. 

The  language  in  which  chemists  express  their  ideas  regard- 
ing the  chemical  constitution  of  bodies  comprises  certain  sym- 
bols or  abbreviations,  to  which  a  conventional  meaning  is 
attached,  and  formulas,  which  are  made  by  grouping  symbols 
together. 

Symbols*  The  symbols  which  stand  in  the  table  after 
the  names  of  the  elements  are  abbreviations,  which  are  used 
instead  of  the  names  in  writing  chemical  formulas. 

Chemical  Formulas  describe  by  means  of  symbols 
the  chemical  constitution  of  bodies. 

Example : — HC1  is  the  formula  of  hydrochloric  acid,  and 
signifies  that  it  is  composed  of  hydrogen  and  chlorine. 

The  Combining  Weights  are  the  numbers  standing 
in  the  table  after  the  symbols.  They  are  peculiar  to  each 
element,  and  denote  the  proportion  by  weight  in  which  it 
unites  with  other  elements.  In  formulas  the  symbols  stand 
for  these  weights,  as  well  as  for  the  names  of  the  elements  ; 
thus,  in  the  formula  of  hydrochloric  acid,  HC1  signifies  that  i 
part  by  weight  of  hydrogen  is  combined  with  35.5  parts  by 
weight  of  chlorine. 

ELEMENTS  COMBINE  WITH  EACH  OTHER  IN  NO  OTHER 

PROPORTIONS    BY    WEIGHT    THAN    THOSE    EXPRESSED    BY    THE 


4  PART  /. 

COMBINING  NUMBERS  (OR  BY  VERY  SIMPLE  MULTIPLES  OF 
THEM). 

This  statement  comprises  the  law  of  definite  proportions 
and  also  that  of  multiple  proportions. 

USUALLY  THE  MULTIPLES  OF  THE  COMBINING  NUMBERS,  2, 
3,  4,  5,  6,  7,  EXPRESS  THE  PROPORTIONS  IN  WHICH  THE  ELE- 
MENTS COMBINE  WITH  EACH  OTHER. 

Example  : — Nitrogen  combines  with  oxygen  only  in  the  pro- 
portions expressed  by  the  formulas  :  N2O,  NO,  N2O3,  NO2, 
and  N2O6,  or,  referring  to  the  table  for  the  combining  num- 
bers for  which  the  symbols  stand,  the  following  proportions 
appear  : 

2  x  14  parts  N  unite  with  16  parts  O. 

14  "  "  "  "  16  "  " 

2  x  14  "  "  "  "  3  x  16  "  " 

14  "  "  "  "  2  x  16  "  " 

2  x  14  "  "  "  "  5  x  16  "  " 

Chemical  Atoms*  It  is  supposed  that  the  utmost  limit 
to  which  the  division  of  matter  could  be  carried  would  lead 
to  its  separation  into  a  great  number  of  particles,  so  small  as 
to  be  incapable  of  further  division.  With  reference  to  their 
quality  of  indivisibility,  such  particles  of  matter  are  called 
atoms  (from  d,  privative,  and  refAvoD,  I  cut).  Atoms,  there- 
fore, are  the  indivisible  constituents  of  matter.  It  is  further 
supposed  that  chemical  combination  consists  in  the  union  of 
atoms,  or  groups  of  atoms,  and  chemical  decomposition  in  the 
separation  of  atoms,  or  groups  of  atoms ;  and  a  chemical 
change  supposes  a  change  in  the  arrangement  or  grouping  of 
the  atoms  of  a  body  involving  the  destruction  of  the  previous 
arrangement. 

Atomic  Weights.  The  atomic  theory  attaches  a  new 
meaning  to  the  combining  weights  of  the  elements,  and  defines 
them  as  the  relative  weights  of  atoms.  Thus,  if  the  weight  of 
an  atom  of  hydrogen  is  i  x  w  (w  =  a  very  small  unknown 


NOTATION.  5 

quantity),  the  weight  of  an  atom  of  chlorine  is  35.5  x  w. 
Not  the  absolute  weights  of  atoms,  but  their  relative  weights, 
have  been  discovered.  The  weight  of  an  atom  of  hydrogen  is 
taken  as  a  standard,  and  called  i  ;  hence  the  weight  of  the 
atoms  of  other  elements  are  expressed  in  terms  of  this  unit. 
Example,  35.5  for  chlorine,  16  for  oxygen.  With  reference 
to  the  above  theory,  the  combining  weights  are  usually  called 
atomic  weights.  The  full  meaning,  therefore,  of  the  chemical 
formula  HC1  is,  that  the  body  which  it  represents  consists  of 
compound  atoms,  each  one  containing  an  atom  of  hydrogen 
and  an  atom  of  chlorine.  An  atom  of  chlorine  weighs  35.5 
times  as  much  as  an  atom  of  hydrogen,  so  that  the  proportion 
by  weight  of  each  constituent  of  the  body  is  expressed  by  its 
formula. 

Chemical  Notation.  Any  change  in  the  constitution 
of  bodies,  as  well  as  their  formation  and  decomposition,  in- 
volves what  is  called  a  chemical  reaction. 

A  CHEMICAL  REACTION  may  be  described  as  a  change  in  the 
arrangement,  or  the  state  of  combination  of  the  atoms  of  bodies. 
Such  a  change  can  be  denoted  by  combining  formulas  together 
in  the  same  way  that  quantities  are  combined  in  common  alge- 
braic calculations.  The  signs  used  are  +,  — ,  and  ='.  Coeffi- 
cients are  only  used  to  multiply  the  symbols  to  which  they  are 
joined.  When  placed  on  the  line,  they  multiply  all  the  sym- 
bols which  follow.  When  placed  below  or  above  the  line, 
they  are  used  to  multiply  only  the  symbol,  or  the  group  of 
symbols  in  brackets,  that  immediately  precedes  them.  Brackets 
are  used  to  distinguish  certain  groups  of  atoms  in  a  compound 
from  the  remaining  atoms.  The  combination  of  atoms  is  ex- 
pressed by  writing  their  symbols  side  by  side,  or  by  grouping 
them  together  without  +  or  — .  The  sign  +  expresses  that 
the  bodies  connected  by  it  are  brought  in  contact  with  each 
other  by  addition,  but  that  they  are  not  combined. 

Example  :— Ba(NO8)2  +  CaSO4.  The  formula  indicates 
that  a  compound,  Ba(NO8)2,  containing  barium,  nitrogen,  and 


6  PARTI. 

oxygen,  in  the  proportion  of  i  atom  of  barium  combined  with 
twice  i  atom  of  nitrogen  and  twice  3  atoms  of  oxygen,  is 
brought  in  contact  with  a  compound  [CaSOJ  containing  cal- 
cium, sulphur,  and  oxygen,  in  the  proportion  of  i  atom  of  cal- 
cium combined  with  i  atom  of  sulphur  and  4  atoms  of  oxygen. 
A  reaction  is  denoted  by  combining  the  formulas  of  the  bodies 
which  take  part  in  it,  as  in  an  ordinary  equation.  The  formulas 
before  the  sign  (  =  )  indicate  the  state  of  combination  of  the 
atoms  before  the  reaction  ;  those  after  the  sign  of  equality  (=) 
show  the  state  of  combination  of  the  atoms  after  the  reaction. 

Example  :— Ba(NO3)a  +  CaSO4  =  BaSO4  +  Ca(NO3)2. 

The  equation  expresses  the  result  of  bringing  in  contact 
the  bodies  described  above,  viz.  :  the  formation  of  new  com- 
pounds containing  barium,  sulphur,  and  oxygen,  and  calcium, 
nitrogen,  and  oxygen. 


CHEMICAL  AFFINITY. 

The  force  which  impels  atoms  to  unite  with  other  atoms  is 
called  chemical  affinity.  The  quantity  or  the  nature  of  the 
force  inherent  in  the  atoms  of  every  substance  determines 
the  chemical  properties  of  the  substance.  The  study  of  the 
results  of  the  action  of  chemical  affinity  is  the  province  of 
chemistry. 

The  phenomena  which  the  action  of  chemical  affinity  gives 
rise  to  can  best  be  studied  under  several  heads. 

Firstly. — Chemical  affinity  may  cause  the  atoms  of  an  ele- 
mentary body  to  unite  among  themselves.  Only  the  cases  of 
such  action  in  which  the  element  is  capable  of  assuming  the 
simplest  physical  condition  of  matter,  namely,  the  form  of  a 
gas,  have  been  studied  satisfactorily. 

The  following  conclusions  in  regard  to  the  state  of  com- 


CHEMICAL  AFFINITY.  7 

bination  of  the  atoms  of  elementary  bodies  have  been  ar- 
rived at : 

The  atoms  of  mercury  and  zinc  remain  separate  in  the 
gaseous  state. 

The  atoms  of  hydrogen,  oxygen,  chlorine,  bromine,  iodine, 
nitrogen  (and  sulphur  at  a  temperature  higher  than  1000° 
centigrade)  are  united  in  groups  of  two  atoms. 

The  atoms  of  phosphorus  and  arsenic  in  the  gaseous  state 
unite  in  groups  of  four  atoms.  Sulphur,  at  a  temperature  of 
500°  centigrade,  in  groups  of  six  atoms. 

The  formulas  for  these  bodies  in  a  gaseous  state  are  : 

Hg        HHorH2  PPPP  or  P4  SSSSSS  or  S6 

Zn          OO  or  O2  AsAsAsAs  or  As4 

C1C1  or  C12 

BrBr  or  Br2 

NN  or  N2 
SS  or  S2 

Molecule.  At  this  point  a  definition  of  the  term  mole- 
cule is  required.  A  molecule  is  the  smallest  particle  of  a  body 
which  can  exist  alone. 

The  molecules  of  elementary  bodies  contain  one  or  more 
atoms  of  the  same  kind.  The  molecules  of  compound  bodies 
contain  two  or  more  atoms  of  different  kinds. 

A  molecule  of  mercury,  hydrogen,  hydrochloric  acid,  or 
water  is  represented  by  the  respective  formulas,  Hg,  H2,  HC1, 
or  H2O. 

Atom.  An  atom  may  be  further  defined  as  the  smallest 
particle  of  matter  which  can  take  part  in  a  chemical  reaction. 
Atoms,  therefore,  appear  while  a  chemical  reaction  is  going 
on,  although  it  is  impossible  to  suppose  that  a  physical  sub- 
division of  matter  could  be  carried  further  than  the  isolation 
of  molecules. 

Thus  in  the  reaction  :  NasS  +  CuCl2  =  CuS  -f  2  NaCl,  the 
force  of  chemical  affinity  breaks  up  the  molecules  (Na^)  and 


8  PART  I. 

(CuCl2)  to  form  the  new  ones  (CuS)  and  (NaCl),  and  during 
this  reaction  the  atoms  Na,  S,  Cu,  and  Cl  must  be  set  free 
from  their  combinations,  and  therefore  must  exist  as  atoms. 

Secondly. — Chemical  affinity  may  combine  atoms  of  a  single 
element,  or  groups  containing  atoms  of  one  or  more  elements, 
with  the  atoms  of  another  element,  or  with  groups  containing 
atoms  of  one  or  more  other  elements. 

Examples : — Hg  -f  C12  =  HgCl2. 

NH3  +  HC1  =  NH4C1. 

The  inverse  action  frequently  takes  place  through  the 
agency  of  heat  or  of  some  other  force,  and  groups  of  atoms 
(molecules)  break  up  into  other  groups,  which  are  usually 
simpler  in  constitution  than  the  primitive  ones. 

Example  :— Hg(CN)2  when  heated  becomes  Hg  +  (CN)2. 

Thirdly. — Chemical  affinity  may  cause  compound  bodies, 
brought  in  contact  with  each  other,  to  mutually  exchange 
some  of  their  constituents ;  or  an  atom  or  a  group  of  atoms 
may  substitute  itself  for  another  atom  or  for  a  group  of  atoms 
in  a  compound  body. 

Examples  :— Na2CO8  +  BaCl2  =  BaCO8  +  2NaCl. 
CuSO4  +  Zn  =  ZnSO4  +  Cu. 

THE  QUALITY  OF  THE  CHEMICAL  AFFINITY  inherent  in  the 
atoms  of  each  element  determines  the  part  which  the  element 
will  play  in  the  different  chemical  changes  mentioned  above. 
It  is  usually  necessary  to  study  each  particular  case,  in  order 
to  determine  the  exact  result  of  bringing  in  contact  any  two 
substances.  Empyrical  rules,  however,  defining  the  nature  of 
the  chemical  affinity  of  the  elements  and  the  consequences  of 
its  action,  can  be  given  in  a  few  cases.  These  rules  are  not 
capable  of  a  very  strict  application,  but  they  serve  to  indicate, 
in  most  cases,  when  a  number  of  bodies  are  brought  together 
in  a  reaction,  those  which  will  probably  combine  with  each 
other.  Gold  is  attacked  by  acids  less  readily  than  the  metals 


CHEMICAL  AFFINITY.  9 

of  the  arsenic  group.     The  difference  between  the  affinities  of 
the  metals  which  are  ranged  in  the  same  group  in  the  table 
(page  2)  is  too  slight  to  be  of  consequence  in  the  application 
of  rule  second. 
First. — COMBINATION  USUALLY  OCCURS  BETWEEN  METALS 

AND  NON-METALLIC  ELEMENTS,  LESS  READILY  BETWEEN  DIF- 
FERENT NON-METALLIC  ELEMENTS,  AND  LEAST  READILY  BE- 
TWEEN DIFFERENT  METALS.  See  table  of  the  elements,  page  2. 

The  following  non-metallic  elements,  or  oxygen  compounds 
of  non-metallic  elements,  combined  with  hydrogen  form  acids, 
and  combined  with  metals  form  salts.  STRONG  ACIDS — sul- 
phuric, H2SO4 ;  nitric,  HNO3  ;  chloric,  HC1O3 ;  chlorhydric, 
HC1.  WEAK  -  ACIDS — sulphurous,  H2SO3 ;  chromic,  H2CrO4 ; 
phosphoric,  H3PO4 ;  boracic,  H3BO3 ;  oxalic,  H2C2O4 ;  acetic, 
HC2H8O2 ;  fluorhydric,  HF  ;  sulphydric,  H2S  ;  cyanhydric, 
HCN  ;  carbonic,  H2CO3 ;  silicic,  H4SiO4.  A  metal  has  a  ten- 
dency to  substitute  itself  for  the  hydrogen  in  a  strong  acid,  to 
form  a  salt  with  it,  in  preference  to  a  weak  one  ;  so  that  a 
strong  acid  usually  displaces  a  weak  one  from  its  salts. 

Second. — THE  METALS  WHICH  STAND  LOWEST  IN  THE  TABLE 

(page  2)  HAVE  THE  GREATEST  TENDENCY  TO  COMBINE  WITH 
THE  STRONGEST  ACIDS. 

Example  : — CuSO*  4-  Zn  =  ZnSO4  +  Cu. 

Third. — Another  rule  which  sometimes  takes  precedence  of 
the  second  is,  that  WHEN  FROM  SOME  OF  THE  CONSTITUENTS 

OF  DIFFERENT  COMPOUNDS  IN  SOLUTION,  AN  INSOLUBLE  BODY 
CAN  BE  FORMED,  THE  ELEMENTS  WHICH  WOULD  COMPOSE 
SUCH  A  BODY  GENERALLY  UNITE  WITH  EACH  OTHER. 

Example:— Na2CO3  +  Ca(HO)2  =  2NaHO  +  CaCO8.  The 
compound  CaCO3  is  formed  in  virtue  of  its  insolubility  when 
the  reaction  takes  place  in  an  aqueous  solution. 


I0  PART  I. 


QUANTIVALENCE. 

The  quality  of  the  chemical  affinity  of  the  elements  deter- 
mines the  nature  of  reactions.  The  quantity  of  their  chemical 
affinity  determines  the  proportions  in  which  the  elements  com- 
bine with  each  other.  The  result  of  the  combination  of  two 
or  more  atoms  with  each  other  is  the  neutralization  of  the 
chemical  force  which  brings  them  together,  so  that  usually  no 
force  is  left  in  the  compound  tending  to  combine  other  bodies 
with  it.  Thus  a  molecule  of  HC1  has  no  power  to  combine 
further  with  atoms  of  H  or  of  Cl.  The  atoms  of  some  ele- 
ments are  animated  with  greater  quantities  of  chemical  force 
than  those  of  other  elements.  Thus  an  atom  of  oxygen  may 
unite  with  an  atom  of  hydrogen,  and  still  be  capable  of  com- 
bining with  another  atom  of  hydrogen,  or  with  an  atom  of 
chlorine.  If  an  atom  of  hydrogen  has  one  unit  of  chemical 
force,  an  atom  of  oxygen  has  two  units,  carbon  has  four,  and 
nitrogen  has  five.  The  number  of  units  of  chemical  force  re- 
siding in  an  atom  is  called  its  QUANTIVALENCE.  The  table 
of  the  elements,  page  2,  classifies  them,  according  to  their 
quantivalence,  into  monatomic  elements,  or  monads  ;  and 
diatomic  elements,  or  dyads,  etc. 

Some  of  the  units  of  chemical  force  of  an  element  may  lie 
dormant  until  developed  by  the  approach  of  some  other  force, 
which  awakens  them.  Iron,  for  instance,  acts  as  a  dyad  when 
no  element  is  present  to  call  forth  all  of  its  four  units  of 
chemical  force.  It  is  worthy  of  note  that  in  such  cases  two 
units  of  force  generally  disappear  together,  as  if  they  became 
dormant  by  each  one  neutralizing  the  effect  of  the  other.  It 
will  be  noticed  that  in  the  table,  page  2,  sulphur  is  placed  in 
the  column  of  dyads,  and  also  in  that  of  the  hexads  ;  this  is 
because  in  many  cases  four  of  the  six  units  of  chemical  force 
in  sulphur  lie  dormant,  and  the  element  plays  the  part  of  a 
dyad.  In  nickel,  cobalt,  and  iron,  two  of  the  four  units  of 


Q  UANTIVA  LENCE.  !  x 

chemical  force  frequently  lie  dormant,  and  hence  these  metals 
often  act  as  dyads.  Only  the  valence  of  each  element,  which 
is  usually  displayed  in  the  kind  of  reactions  which  we  have  to 
consider,  is  shown  in  the  table. 

The  knowledge  of  these  facts  is  essential  to  enable  the  stu- 
dent to  write  formulas  correctly. 

Examples: — The  compound  of  barium  and  chlorine  must 
contain  two  atoms  of  chlorine,  combined  with  one  atom  of 
barium.  Its  formula  is  BaCl2.  The  greatest  amount  of  oxy- 
gen that  an  atom  of  carbon  can  unite  with  is  two  atoms. 
The  formula  of  the  compound  is  CO2.  When  zinc  is  substi- 
tuted for  silver  in  a  compound,  one  atom  of  zinc,  with  its  two 
units  of  chemical  force  or  affinity,  takes  the  place  of  two 
atoms  of  silver,  because  each  atom  of  silver  has  only  one  unit 
of  chemical  force.  Thus,  in  writing  formulas,  an  atom  of  one 
element  is  equivalent  to  or  takes  the  place  of  another  element 
of  the  same  class.  In  comparing  elements  of  different  classes, 
their  value  in  an  equation  depends  upon  the  number  of  units 
of  chemical  force  which  they  contain. 

TWO  MONADS  ARE  EQUIVALENT  TO  A  DYAD  ;  THREE  MO- 
NADS TO  A  TRIAD,  ETC.  THREE  DYADS  ARE  EQUIVALENT  TO 
TWO  TRIADS,  ETC. 

Examples :— PbO  +  2HC1  =  PbCl2  +  H2O. 
SiCl4  +  2H2O  =  SiO8  +  4HC1. 

QUANTIVALENCE  OF  GROUPS  OF  ATOMS. — When  part  of  the 
affinities  or  units  of  chemical  force  of  a  polyatomic  group  are 
satisfied  or  neutralized,  the  residual  valences,  or  those  which 
remain  free,  determine  the  quantivalence  of  the  group. 

Example : — Nitrogen,  a  pentad,  when  combined  with  three 
atoms  of  hydrogen,  is  capable  of  uniting  with  two  other  mo- 
nads or  with  one  dyad — NH3  can  unite  with  H  and  Cl  to  make 
NH4C1.  Moreover,  certain  groups  of  elements  take  part  in 
many  chemical  reactions  without  being  broken  up,  and,  so  far 
as  this  is  the  case,  they  may  be  considered  as  playing  the  part 


12  PART  L 

of  elementary  bodies  in  the  reactions.  Such  groups  are  fre- 
quently called  radicals,  with  reference  to  a  theory  that  they 
are  the  roots  of  compounds. 

The  following  table,  showing  the  quantivalence  of  such 
groups,  will  be  found  convenient  in  writing  formulas : 

Monads.  Dyads.  Triads.  Tetrads.  Hexads. 

HO*  SO4  PO4  SiO4  Fe2§ 

N08f  S03  B03  Fe(CN)6     Cr2 

C103  Cr04  A12 

C2H3O2  CO3  Fe2(CN)12 

NH4  C2O4 

CN  C4H4O6 
Hg2 

Examples :— Na(HO);  H(NO3);  (NH4)  (HO);  Pb(C2H8O2)2; 

*  The  group  HO  is  a  monad,  because  one  of  the  two  affinities  of  the 
oxygen  atom  is  satisfied  and  nullified  by  the  affinity  of  the  hydrogen  with 
which  it  is  combined,  leaving  one  affinity  free  in  the  oxygen  atom. 

f  The  existence  of  the  group  NO3  might  seem  inconsistent  with  the 
five-atomic  character  of  nitrogen.  The  explanation  is,  that  four  affinities 
in  the  nitrogen  atom  are  used  to  unite  to  it  two  atoms  of  oxygen,  and  the 
fifth  combines  with  only  one  of  the  affinities  of  a  third  atom  of  oxygen, 
leaving  one  oxygen  affinity  free,  which  determines  the  monatomic  char- 
acter of  the  group. 

\  A  similar  argument  applies  to  the  case  of  the  group  SO4.  Here,  sul- 
phur being  hexatomic,  four  affinities  are  used  to  combine  with  those  of  two 
atoms  of  oxygen,  while  the  two  remaining  affinities  of  the  sulphur  atom 
are  each  combined  with  one  in  each  of  the  two  remaining  atoms  of  oxy- 
gen, leaving  two  oxygen-affinities  free. 

§In  the  case  of  Fe2,  etc.,  two  of  the  eight  affinities,  belonging  to  two 
atoms  of  iron,  are  used  to  bind  the  two  atoms  together,  leaving  six  free, 
and  for  this  reason  the  group  Fe2  is  six-,  and  not  eight-atomic.  It  may 
be  noticed  that  groups  of  this  nature,  having  free  chemical  affinities,  can 
only  be  found  in  the  case  of  polyatomic  elements,  because,  when  two 
monatomic  elements  combine  with  each  other,  no  chemical  affinity  can 
be  left  free.  For  fuller  explanations  of  the  laws  of  chemical  combination, 
see  Frankland's  Lecture  Notes  for  Chemical  Students. 


QUANTIVALENCE.  !3 

H2(S04)  ;   Ba  (CrO4) ;  Hg2  (C12)  ;  H3  (PO4)  ;  Fe2Cl6 ;  Cr2O, ; 
A12  (HO).. 

IN  WRITING  FORMULAS,  THE  SYMBOL  OF  THE  ELEMENT,  OR 
GROUP  OF  ELEMENTS,  HAVING  MOST  DECIDEDLY  THE  CHARAC- 
TER OF  A  METAL,  IS  PLACED  FIRST. 

Example :— NaNH4HPO«. 


I4 


CHAPTER  II. 

CHEMICAL   NOMENCLATURE. 

WHEN  metallic  sodium  comes  in  contact  with  water  a  reac- 
tion takes  place,  as  indicated  by  the  following  equation  : 

Na*  +  2H2O  =  2NaHO  +  H2. 

Each  molecule  of  water  is  broken  up,  setting  free  an  atom 
of  hydrogen  and  receiving  in  its  place  an  atom  of  sodium. 
The  new  body  thus  formed  may  then  be  considered  as  a  mole- 
cule of  water  in  which  one  atom  of  hydrogen  has  been  replaced 
by  sodium.  Or  it  may  be  considered  as  an  atom  of  a  monad 
metal  united  to  the  monovalent  radical  HO,  known  as  hy- 
droxyl.  Furthermore,  it  may  be  taken  as  a  type  of  a  numerous 
class  of  bodies,  known  as  METALLIC  HYDRATES,  all  of  which 
consist  of  one  or  more  atoms  of  metal  united  to  an  equivalent 
amount  of  hydroxyl,  as  may  be  seen  from  the  following 

Examples:— KHO  ;  Ca(HO)2 ;  Fe(HO)2 ;  Fe2  (HO),  ;  Al, 
(HO), 

All  of  the  metallic  hydrates,  when  acted  upon  by  a  certain 
other  class  of  bodies,  are  similarly  affected,  as  can  be  shown 
by  the  following  equations  : 

HC1      +  KHO         =  KC1  +     H20 ; 

2HC1      +  Ca(HO)2   =  CaCl2         +  2H2O  ; 

6HC1      +  Fe2(HO)6  =  Fe2Cl6        +  6H2O  ; 

HNO8  +  KHO         =  KNO3        +     H2O  ; 

6HN08  +  Fe2(HO)6  =  Fe2(NO8)6  +  6HO ; 

H2S04  +  KHO        =  KHS04     +    H2O ; 

H2SO4  +  2KHO       =  K2SO4        +  2H2O  ; 

H2SO<  +  Ca(HO)2  =  CaS04       +  2H2O. 


CHEMICAL  NOMENCLATURE.  !5 

By  inspection  of  the  above  equations  it  will  be  seen  that  we 
have  in  the  left-hand  member  of  each  a  metallic  hydrate,  and 
in  the  right-hand  member  of  each  water.  Furthermore  the 
substance  indicated  first  in  each  equation  contains  hydrogen, 
and  when  this  hydrogen  is  replaced  by  the  metal  of  the  metal- 
lic hydrate  the  formula  of  the  substance  indicated  first  in  the 
right-hand  member  of  the  equation  is  obtained. 

The  original  substances  containing  hydrogen  are  called 
acids,  and  the  compounds  occurring  first  in  the  right-hand 
members  of  the  above  equations  are  salts. 

The  above  we  may  formulate,  then,  in  the  following  defi- 
nitions : 

Acid.  An  acid  is  a  compound  which,  when  brought  into 
contact  with  a  metallic  hydrate,  exchanges  the  whole  or  a  part 
of  its  hydrogen  for  the  metal  of  the  metallic  hydrate ;  water 
being  formed  at  the  same  time. 

The  acids  may  be  divided  into  two  groups. 

Firstly,  Those  which  consist  of  hydrogen  united  directly  to 
another  non-metallic  element  as  HC1  ;  HBr  ;  HF  ;  H2S. 

Secondly,  The  oxygen  acids,  or  those  in  which  the  hydrogen 
is  united  with  another  non-metal,  or  with  a  group,  by  means  of 
oxygen,  as  in  HC1O  ;  HNOa  ;  H2SO4.  Here  the  arrangement 
of  the  chemical  forces  may  be  shown  by  the  following  graphic 
formulas  :  H  -  O  -  Cl  ;  H  -  O  -  N  O2 ;  H2  =  O2  --=  SO2. 

Salt .  A  salt  is  a  compound  formed  by  replacing  the  hy- 
drogen of  an  acid  by  a  metal. 

All  compound  bodies  are  susceptible  of  decomposition  by  a 
current  of  electricity,  and  when  thus  decomposed  resolve  them- 
selves into  a  positive  atom  or  group,  and  a  negative  atom  or 
group.  In  writing  the  ordinary  formulas  of  compounds  the 
electro-positive  atom  or  group  is  always  placed  first,  as 

+  -      +  -      +    -         +     -  + 

HC1  ;    H20  ;  Na(HO)  ;  Ba(SO4)  ;  (NH4)2(SO4). 

The  names  of  chemical  compounds,  which  are  used  in  works 
on  analytical  chemistry,  do  not  describe  their  constitution  as 


16  PART  I. 

fully  as  formulas  do  ;  they  serve,  however,  to  identify  the 
bodies  and  to  recall  certain  principles  of  classification. 

In  the  ordinary  operations  of  analysis,  the  only  bodies  which' 
are  dealt  with  are  binary  compounds  (i.e.,  compounds  which 
only  contain  two  elements,  as  KC1  and  H2O)  and  ternary  com- 
pounds, which  contain  a  non-metallic  element  and  oxygen 
combined  with  a  metal  or  with  hydrogen,  as  BaSO4 ,  H3PO4. 

In  the  statement  of  the  rules  for  the  naming  of  compounds 
the  subject  will  be  divided  under  these  two  heads  : 

I.  BINARY  COMPOUNDS.       * 

a.  ACIDS. — Two  distinct  methods  are  in  common  use  for 
naming  binary  acids.     According  to  the  one  the  name  consists 
of  hydro  preceding  the  root  of  the  name  of  the  other  element 
and  terminates  in  ic;  as,  HC1,  hydrochloric  acid  ;  HBr,  hydro- 
bromic  acid  ;  H2S,  hydrosulphuric  acid.     By  the  other  method 
the  characteristic  element  is  indicated  first,  and  the  names  of 
the  three  acids  just  mentioned  would  be,  chlorhydric  acid, 
bromhydric  acid,  and  sulphydric  acid. 

b.  ALL  OTHER  COMPOUNDS. — In  determining  the  names  of 
binary  compounds  not  acids,  if  there  exists  but  one  compound 
of  the  two  elements  in  question,  the  positive  element  takes  the 
termination  ic  and  the  negative  ide;  as,  NaCl,  sodic  chloride  ; 
K2S,  potassic  sulphide  ;  BaO,  baric  oxide.     Or,  inverting  the 
order,  they  may  also  be  called,  respectively,  chloride  of  sodium, 
sulphide  of  potassium,  and  oxide  of  barium. 

Still  another  method,  and  one  which  has  of  late  been  very 
generally  adopted,  is  to  leave  the  name  of  the  positive  element 
unchanged,  while  giving  the  termination  ide  to  the  negative. 
By  this  method  the  examples  just  given  would  be  sodium 
chloride,  potassium  sulphide,  and  barium  oxide. 

When  two  distinct  compounds  exist,  consisting  of  the  same 
two  elements,  as  SO2,  SO3,  that  one  which  contains  the  greater 
proportion  of  the  negative  element  is  named  in  accordance 


CHEMICAL  NOMENCLATURE.  ^ 

with  the  first  method  as  given  above,  while  the  other  simply 
changes  the  termination  ic  of  the  positive  element  to  ous.  SO2 
would  then  be  sulphurous  oxide  and  SO3  sulphuric  oxide. 

When  a  compound  exists  which  contains  still  less  of  the 
negative  constituent  than  that  indicated  by  the  termination 
ousy  as  above,  it  receives  the  prefix  hypo ;  as,  N2O3  being  ni- 
trous oxide,  N2O  is  called  hyponitrous  oxide.  Similarly  the 
prefix  per  or  hyper  is  used  to  indicate  that  the  compound  con- 
tains more  of  the  negative  element  than  the  one  whose  positive 
constituent  terminates  in  ic.  This  prefix  is,  however,  usually 
given  to  the  negative  element ;  as,  Mn2O3  being  manganic  oxide, 
MnO2  is  called  manganic  peroxide. 

The  prefix  sesqui  is  sometimes  used,  as  in  Fe2O3,  the  sesqui 
oxide  of  iron. 

The  relative  positions  of  the  several  compounds  of  a  series 
may  be  illustrated  by  the  following  schedule,  in  which  the  plus 
sign  indicates  the  root  of  the  positive  element  and  the  minus 
sign  that  of  the  negative : 

hypo  +  ous         ide,  as  XO  ; 

4-  ous         ide,  as  XO2 ; 

hypo  +  ic  ide,  as  XO3 ; 

+  ic  ide,  as  XO4 ; 

+  ic  per ide,  as  XO5. 

Unfortunately  for  the  rule,  however,  the  practice  of  recent 
authors  varies  very  widely,  so  much  so  that  it  is  the  exception, 
for  example,  to  find  the  whole  series  of  oxides  of  nitrogen 
similarly  named  in  any  two  works  on  chemistry. 

In  accordance  with  the  above  schedule  the  following  would 
be  the  names  of  the  five  members  of  that  series  : 

N2O,   hyponitrous  oxide  ; 
NO,    nitrous  oxide  ; 
N2O3,  hyponitric  oxide ; 
NO2,   nitric  oxide ; 
N2O5,  nitric  peroxide. 


l8  PART  I. 

Instead  of  these,  however,  each  member  of  the  series  has 
several  different  names  in  common  use.  For  example,  the  first 
compound  is  known  commonly  as  nitrous  oxide  ;  it  is  also 
called  nitrogen  monoxide.  The  second  is  variously  termed 
nitric  oxide,  nitrosyl,  nitrogen  dioxide,  since  the  molecule  is 
by  some  considered  to  be  indicated  by  the  formula  N2O2,  and 
so  on  through  the  whole  of  the  series. 

Another  striking  example  of  exception  to  the  rule  will  be 
found  in  the  two  compounds  of  carbon  and  oxygen,  CO  and 
CO2.  Clearly  the  former  should  be  called  carbonous  oxide, 
and  the  latter  carbonic  oxide.  Formerly,  however,  CO2  was 
considered  to  be  an  acid,  and  was  called  carbonic  acid  ;  and 
in  fact  it  is  still  so  termed  in  most  metallurgical  and  technical 
works,  while  CO  was  known  as  carbonic  oxide.  Since  the 
adoption  of  the  present  views  of  the  constitution  of  acids 
chemists  no  longer  call  CO2  carbonic  acid,  and  it  is  therefore 
usually  named  carbonic  dioxide,  or  carbonic  anhydride,  from 
the  fact  that  it  may  be  considered  as  carbonic  acid  from  which 
a  molecule  of  water  has  been  removed  (HaCO8  —  H2O  =  CO2). 
The  name  carbonic  oxide  would  be  inapplicable,  since  it  would 
not  be  perfectly  clear  whether  the  compound  indicated  was 
CO  or  CO2.  As  to  CO,  it  still  retains  its  old  name  carbonic 
oxide  in  some  works,  while  in  others  it  is  termed  carbonous  ox- 
ide, or  again  carbon  monoxide. 

Still  another  source  of  difficulty  to  beginners  exists  in  the 
fact  that  many  compounds  are  known  by  irregular  names ;  as, 
H2O,  water ;  NaCl,  common  salt  ;  NH8,  ammonia,  etc. 

For  the  sake  of  uniformity  it  would  perhaps  be  well  if  the 
system  as  shown  in  the  following  examples  could  be  adopted : 

N2O,   nitrogen  monoxide'; 
N2O2,  nitrogen  dioxide ; 
N2O3,  nitrogen  trioxide ; 
N2O4,  nitrogen  tetroxide ; 
N2O5,  nitrogen  pentoxide ; 
or,  nitric  monoxide,  etc. 


TERNARY  COMPOUNDS.  !9 

Mn8O4  would,  according  to  this   plan,  be   called  triman- 
ganese  tetroxide. 

II.  TERNARY  COMPOUNDS. 

a.  ACIDS. — The  elements  hydrogen  and  oxygen  are  common 
to  all  of  this  class  of  acids.     They  therefore  receive  their  speci- 
fic names  from  the  third  or  characteristic  element ;  as,  HNO8 
is  called  nitric  acid.     In  case  two  or  more  acids  occur,  consist- 
ing of  the  same  three  elements  united  in  different  proportions, 
the  prefixes  and  terminations,  as  explained  under  binary  com- 
pounds, are  applied,  and  in  a  similar  manner ;  as, 

HNO2,  nitrous  acid  ; 
HNO8,  nitric  acid ; 
H2SO3,  sulphurous  acid ; 
H2SO4,  sulphuric  acid ; 
HC1O,  hypochlorous  acid ; 
HC1O2,  chlorous  acid  ; 
HC1O3,  chloric  acid ; 
HC1O4,  perchloric  acid. 

b.  SALTS. — The  names  of  salts  express  both  the  metal  con- 
tained and  the  acid  from  which  they  are  derived.     The  metal 
takes  the  termination  ous  or  ic.     The  termination  of  the  acid 
is  changed  from  ous  to  tie,  and  from  ic  to  ate  j  as, 

KNO2,  potassic  nitrite ; 
NaNO3,  sodic  nitrate  ; 
BaSO4,  baric  sulphate ; 
FeSO4,  ferrous  sulphate ; 
Fe2(SO4)8,  ferric  sulphate ; 
KC1O,  potassic  hypochlorite ; 
NaClO2,  sodic  chlorite  ; 
KC1O8,  potassic  chlorate  ; 
KC1O4,  potassic  perchlorate. 


20  PART  I. 

c.  METALLIC  HYDRATES. — The  metal  generally  takes  the 
termination  /V,  as  NaHO,  sodic  hydrate.  Where  two  hydrates 
of  the  same  metal  exist  the  one  which  has  the  lesser  relative 
amount  of  the  negative  radical  takes  the  termination  ous;  as, 
Fe(HO)2,  ferrous  hydrate  ;  Fe2(HO)6,  ferric  hydrate. 

In  cases  where  only  one  hydrate  is  formed  some  authors 
prefer  to  retain  the  name  of  the  metal  unchanged  ;  as,  sodium 
hydrate. 

CLASSIFICATION   OF    ACIDS. 

The  hydrogen  of  an  acid  which  is  replaceable  by  a  metal  is 
called  basic  hydrogen. 

Those  acids  which  contain  in  the  molecule  but  one  atom  of 
basic  hydrogen  are  said  to  be  monobasic  ;  those  which  con- 
tain two  such  atoms  are  dibasic  ;  those  containing  three  are 
tribasic,  and  those  containing  four,  tetrabasic.  In  hypophos- 
phorous  acid,  H3PO2,  we  have  an  example  of  a  monobasic 
acid,  since  only  one  of  the  three  atoms  of  hydrogen  is  replace- 
able. In  phosphorous  acid,  H3PO8,  we  have  a  dibasic  acid, 
and  in  phosphoric  acid,  H8PO4,  we  have  a  tribasic  acid,  all 
three  atoms  of  hydrogen  being  basic.  Pyrophosphoric  acid, 
H4P2O7,  is  an  example  of  a  tetrabasic  acid. 

CLASSIFICATION   OF   SALTS. 

Salts  are  either  normal,  acid,  double,  or  basic. 

A  normal  salt  is  one  in  which  the  whole  of  the  basic  hydro- 
gen is  replaced  by  the  metal ;  as,  NaNO3 ;  Na2SO4 ;  NaH2PO2 ; 
Na,2HPO8;  Na8PO4. 

An  acid  salt  is  one  in  which  some  of  the  basic  hydrogen 
still  remains  in  the  molecule  ;  as,  NaHSO4 ;  NaH2PO3  ; 
Na2HPO4.  There  can  of  course  be  no  acid  salt  of  a  mono- 
basic acid,  since,  as  the  molecule  contains  but  one  atom  of  hy- 
drogen, if  any  portion  of  that  element  is  removed  the  whole 
must  be. 


CHEMICAL  REAGENTS.  21 

A  double  salt  is  a  complex  molecule  in  which  either  the 
basic  hydrogen  is  replaced  by  two  distinct  metals,  or  else  the 
same  metal  replaces  the  basic  hydrogen  in  two  different  acids. 
As  examples  of  the  first  variety  of  double  salt  may  be  men- 
tioned (KCl)2PtCl4,  in  which  the  two  metals,  potassium  and 
platinum,  replace  the  hydrogen  in  hydrochloric  acid  ;  CaMg 
(CO3)2 ;  K2A12(SO4)4.  Examples  of  the  second  form, 

PbS04,  3PbC08 ;  PbCl2,  PbCO3. 

Basic  salts  are  complex  molecules  which  may  be  consid- 
ered as  consisting  of  a  normal  salt  united  to  the  oxide  or  the 
hydrate  of  the  same  metal  ;  as,  BiCl3,  Bi2O3  ;  Fe2(SO4)3, 
Fe2(HO)6 ;  (CuCO8)2,  Cu(HO)2. 

It  will  be  seen  by  the  foregoing  examples  how  much  more 
fully  the  formulas  express  the  composition  of  bodies  than  the 
names. 

The  following  list  of  chemical  compounds  contains  the 
names  and  formulas  of  the  substances  most  frequently  used 
for  testing  in  the  laboratory.  These  names  and  formulas 
should  be  committed  to  memory,  and  the  rules  of  nomencla- 
ture may  be  studied  in  their  application  to  them  : 

Chlorhydric,  or  hydrochloric  Potassic  ferrocyanide, 

acid,  HC1  ;  K4(FeCy6),  or  K,Ye(CN)6 ; 

Nitric  acid,  HNO3 ;  Potassic  ferricyanide, 

Sulphuric  acid,  H2SO4 ;  K6(Fe2Cy,2),  or  K6Fe2(CN)12; 

Sulphydric  acid,  H2S  ;  Potassic  sulphocyanate, 

Acetic  acid,  HC2H3O2 ;  K(CyS),  or  KCNS  ; 

Ammonic  hydrate,  NH4HO  ;  Calcic  hydrate,  Ca(HO)2 ; 

Ammonic  sulphide,  (NH4)2S  ;  Calcic  chloride,  CaCl2 ; 

Ammonic  carbonate,  Calcic  sulphate,  CaSO4 ; 

(NH4)2CO3 ;  Baric  chloride,  BaCl2 ; 

Ammonic  chloride,  NH4C1  ;  Baric  nitrate,  Ba(NO3)2 ; 

Ammonic  oxalate,  Baric  carbonate,  BaCO3 ; 

(NH4)2C2O4 ;  Magnesic  sulphate,  MgSO4 ; 


22 


PART  /. 


Ferrous  sulphate,  FeSO4 ; 
Ferric  chloride,  Fe2Cl6 ; 
Cobaltic  nitrate,  Co(NO8)2 ; 
Plumbic  acetate  ; 

Pb(C2H302)2; 
Argentic  nitrate,  AgNO3 ; 
Mercuric  chloride,  HgCl2 ; 
Platinic  chloride,  PtCU  ; 
Stannous  chloride,  SnCl2 ; 
Alcohol,  C2H6O. 

After  having  made  himself  familiar  with  the  formulas  of  the 
substances  with  which  he  tests,  the  student  should  write  in  the 
form  of  an  equation  the  result  of  the  action  of  each  test  which 
he  performs  upon  a  compound  under  examination. 


Ammonic  molybdate, 

(NH4)2Mo04 ; 
Sodic  hydrate,  NaHO  ; 
Sodic  carbonate,  Na2CO8 ; 
Disodic  hydric  phosphate, 

Na2HPO4 ; 

Sodic  acetate,  NaC2H8O2 ; 
Potassic  dichromate, 

K2Cr2O7 ; 


CHEMICAL   OPERATIONS. 


CHAPTER  III. 

CHEMICAL   OPERATIONS. 

Reaction  with  Test  Paper.  A  small  piece  of  red 
or  blue  litmus  paper  is  dipped  in  the  solution.  If  it  is  acid, 
blue  paper  is  turned  red  ;  if  it  is  alkaline,  red  paper  is  turned 
blue. 

Turmeric  paper  is  turned  brown  by  alkalies. 

Precipitation.  When  an  insoluble  body  is  formed  in 
a  solution  and  separates  (falls)  from  it,  precipitation  is  said  to 
take  place.  Precipitates  are  gelatinous,  as  aluminic  hydrate ; 
flocculent  (consisting  of  flakes),  as  sulphide  of  zinc  ;  or  pul- 
verulent, as  baric  sulphate.  Usually  the  particles  are  less  finely 
divided,  and  filtration  is  easier  with  precipitates  which  form  in 
dilute  solutions,  particularly  in  boiling  solutions.  With  some 
precipitates,  as  magnesic  phosphate,  in  very  dilute  solutions, 
the  act  of  precipitation  is  a  slow  process  of  crystallization,  and 
the  formation  of  a  precipitate  does  not  take  place  until  after 
several  hours. 

Filtration  is  a  process  by  which  an  insoluble  body, 
usually  a  precipitate,  is  separated  from  a  liquid.  It  is  usually 
important  to  allow  a  precipitate  to  settle  before  filtration  ;  and 
frequently  after  the  clear  liquid  has  been  poured  upon  the 
filter  it  is  best  to  add  more  water  to  the  precipitate,  and  to 
wait  again  until  it  has  settled.  When  the  precipitate  requires 
to  be  washed,  this  process  may  be  repeated  many  times  before 
the  precipitate  is  brought  upon  the  filter.  The  precipitate 
usually  clogs  the  pores  of  the  filter,  so  that  it  retards  the  flow 
of  the  liquid.  A  precipitate  which  has  to  be  washed  is  finally 
brought  on  the  filter  by  rinsing  with  the  wash-bottle,  and  it  is 


24  PART  I. 

washed  by  repeatedly  filling  the  filter  with  water  and  allowing 
it  to  empty  itself.  The  process  should  be  intermittent,  and 
the  filter  should  never  be  kept  constantly  full  during  the  latter 
part  of  a  filtration,  when  the  object  is  to  wash  a  precipitate 
with  pure  water.  Sometimes  precipitates  which  take  the  form 
of  a  powder  are  so  finely  divided  as  to  pass  through  the  pores 
of  a  filter.  This  can  usually  be  avoided  by  precipitating  in  a 
dilute  solution,  and  particularly  by  boiling  the  solution.  Sul- 
phur cannot  be  prevented  from  going  through  the  filter.  Fil- 
tration is  more  rapid  with  hot  than  with  cold  water.  Before 
commencing  a  filtration,  the  filter  should  be  made  to  fit  closely 
to  the  side  of  the  funnel,  and  it  should  always  be  moistened 
with  pure  water. 

Decatfltatiovi  consists  in  allowing  a  precipitate  to  settle 
and  in  pouring  off  the  liquid  above  it,  in  the  manner  already 
mentioned  under  filtration.  This  process  may  or  may  not  be 
united  with  that  of  filtration.  For  instance,  argentic  chloride 
settles  so  completely  and  quickly  that  it  can  usually  be  washed 
simply  by  repeated  decantations. 

^Evaporation  is  usually  performed  in  a  porcelain  dish. 
A  few  drops  of  liquid  can  be  evaporated  by  heating  them  on 
platinum  foil.  Bits  of  broken  glass  or  porcelain  vessels  are 
very  useful  for  the  same  purpose. 

The  Use  of  the  Blowpipe.  An  olive  oil  or  kerosene 
lamp,  with  a  wick  $  in.  long  and  J  in.  broad,  or  a  Bunsen's 
lamp,  with  the  regulator  turned  to  shut  off  the  draught  of  air, 
may  be  used.  If  the  Bunsen's  lamp  has  no  regulator  for  the 
draught,  a  smaller  tube  may  be  introduced  into  the  lamp-tube, 
until  it  rests  upon  the  piece  from  which  the  gas  issues  and 
excludes  the  air.  After  considerable  practice  a  continuous 
stream  of  air  can  be  forced  through  the  blowpipe  by  making 
the  cavity  of  the  mouth  the  reservoir,  into  which  air  is  forced 
at  intervals  from  the  lungs,  and  is  prevented  from  escaping  by 
a  peculiar  contraction  of  the  throat  and  hanging  palate,  which 
is  easily  learned  ;  the  breathing  goes  on  through  the  nose  un- 


THE    USE   OF   THE  BLOWPIPE.  25 

interruptedly.  The  chief  difficulty  that  beginners  usually  ex- 
perience is  fatigue  of  the  muscles  of  the  cheeks,  which  pre- 
vents a  long-continued  effort.  Two  kinds  of  flames  can  be 
produced  with  the  blowpipe — one  containing  an  excess  of  air, 
consequently  of  oxygen  ;  the  other  containing  an  excess  of 
combustible  gases,  consequently  of  gases  capable  of  consuming 
oxygen  or  reducing. 

The  Oxidizing  Flame  is  produced  by  introducing 
the  point  of  the  blowpipe  one-third  through  the  lamp-flame. 
It  is  a  clear  blue  cone,  surrounded  and  continued  at  the  point 
by  a  colorless  flame,  intensely  hot,  and  capable  of  producing 
oxidation.  The  substance  should  be  heated  beyond  the  point 
of  the  blue  cone. 

TJie  Reducing  Flame  is  produced  by  holding  the 
point  of  the  blowpipe  at  the  outside  of  the  lamp-flame,  and 
by  blowing  somewhat  more  gently.  The  flame  is  much  less 
pointed  and  is  more  luminous  than  the  oxidizing  flame.  The 
substance  should  be  heated  at  a  distance  from  the  point  of  the 
flame  equal  to  one-third  of  its  length,  and  should  be  completely 
enveloped  in  the  flame.  The  position  of  the  blowpipe  and 
the  force  of  the  blast  regulate  the  quality  of  the  flame.  It  is 
a  difficult  matter  to  produce  the  true  reducing  flame,  which 
should  not  deposit  carbon  on  the  substance  heated  in  it,  and 
at  the  same  time  should  contain  no  excess  of  oxygen. 

The  Borax  Bead  is  formed  by  making  a  loop  -J-  in. 
in  diameter  in  a  piece  of  platinum  wire,  heating  it  red  hot  in 
the  blowpipe  flame,  and  touching  it  to  a  small  piece  of  borax 
while  it  is  hot.  The  borax,  which  adheres  to  the  hot  wire,  is 
heated  in  the  blowpipe  flame.  It  at  first  swells  while  losing 
its  water  of  crystallization,  and  finally  it  melts  to  a  clear  glass 
bead.  A  finely  divided  substance  can  be  taken  up  by  touch- 
ing the  hot  bead  to  it,  and  it  can  then  be  tested  as  to  its  solu- 
bility in  the  borax  bead,  coloring  properties,  etc.,  in  the  blow- 
pipe flame. 

Heating  on  Charcoal.    Select  a  good  piece  of  char- 


26  PART  I. 

coal,  at  least  4  in.  long  and  i  in.  broad  and  thick,  and  smooth 
a  plane  surface  in  a  direction  at  a  right  angle  with  that  of  the 
year-rings.  (If  heat  is  applied  to  a  surface  parallel  to  the 
planes  of  the  year-rings,  the  charcoal  is  more  liable  to  snap 
from  the  expulsion  of  moisture.)  In  many  cases  the  charcoal 
serves  as  a  convenient  support  for  a  substance  to  be  heated  ; 
in  others  the  reducing  agency  of  the  charcoal  comes  in  play. 
Substances  are  also  evaporated  at  a  high  temperature  from  the 
surface  of  the  charcoal. 

Ductility,  Malleability,  Urittleness  are  charac- 
teristic properties  of  metals,  and  metals  can  be  tested  with  re- 
gard to  them  by  pounding  with  a  hammer  or  by  rubbing  with 
the  pestle  of  a  mortar.  When  a  substance  which  appears  to 
contain  a  higher  metal  is  reduced  by  sodic  carbonate  on  char- 
coal, unless  metallic  globules  are  at  once  apparent,  the  portion 
of  the  charcoal  which  has  been  heated  should  be  cut  out  and 
pulverized  with  water  in  a  mortar,  and  washed  by  decantation. 
If  metallic  globules  have  been  formed,  they  will  sink  to  the 
bottom,  and  after  thorough  washing,  during  which  the  sodic 
carbonate  is  dissolved,  and  the  light  particles  of  charcoal  are 
floated  away,  they  will  appear  as  globules,  if  they  are  hard 
like  copper  ;  as  brittle  grains,  if  they  are  brittle  like  bismuth  ; 
or  as 'flattened  disks,  when  they  are  ductile  like  lead,  and  when 
they  have  been  pressed  by  the  pestle  against  the  mortar. 

Color  of  the  Flame.  If  the  substance  to  be  tested  is 
a  solid,  a  small  piece  of  it  is  brought  on  a  loop  of  fine  platinum 
wire,  or  in  a  pair  of  forceps,  into  the  flame  of  the  alcohol  or 
Bunsen  lamp,  and  the  color  imparted  to  the  flame  observed. 
If  a  substance  in  solution  is  to  be  tested,  the  platinum  wire  is 
dipped  in  the  solution  and  is  then  introduced  into  the  lamp- 
flame.  If  the  solution  is  too  dilute  to  afford  a  distinct  test  in 
this  way,  it  must  be  evaporated,  and  it  is  usually  best  to  evapo- 
rate nearly  to  dryness,  and  to  take  some  of  the  solid  residue 
for  the  test. 

Tlie  Manipulation  of  Glass  Tubing.    Glass 


CHEMICAL   OPERATIONS.  27 

softens  when  a  small  piece  is  heated  in  the  flame  of  an  alcohol 
lamp,  or  when  a  larger  piece  is  heated  in  a  Bunsen  lamp,  or 
with  the  blowpipe  flame  or  in  a  blast  lamp. 

A  tube  can  be  bent  easily  as  soon  as  the  glass  softens. 

It  is  best  only  to  bend  gently  at  first,  then  to  .heat  the  adja- 
cent part  of  the  tube,  and  to  bend  again,  and  so  on,  in  order 
that  the  sides  of  the  tube  may  not  fall  together  in  bending. 

When  a  glass  tube  is  heated  for  some  time,  it  contracts  and 
the  sides  thicken.  By  drawing  out  a  tube  either  immediately 
after  it  has  become  soft,  or  after  the  sides  have  thickened,  a 
tapering  point  of  any  desired  calibre  and  thickness  of  glass  can 
be  obtained. 

To  close  the  end  of  a  glass  tube,  draw  the  tube  off  while  the 
glass  is  as  thin  as  possible,  and  hold  the  tapering  point  in  the 
flame,  and  draw  the  end  off  again  ;  in  this  way  a  tube  with  a 
pointed  closed  end  is  obtained  ;  by  heating  the  closed  end  of 
the  tube,  removing  it  from  the  flame,  rotating  it  and  blowing 
in  it,  while  the  glass  is  still  red  hot,  it  expands,  and  a  more 
rounded  end  or  a  bulb  can  be  produced.  Glass  tubing  can  be 
cut  by  making  a  mark  at  the  required  place  by  a  few  file 
strokes,  and  then  by  breaking  the  glass.  The  ends  of  glass 
tubes  cut  in  this  way  should  be  held  in  the  flame  till  they 
become  red  hot ;  in  this  way  the  sharp  edges  become  rounded. 


PART  II. 


PART  II.  is  preparatory  to  Part  III.,  which  contains  a  general 
scheme  of  analysis  applicable  to  compounds  of  all  the  elements. 
The  most  important  tests  are  those  which  are  described  in  Part 
III.,  and  a  knowledge  of  them  would  suffice  alone  for  the  pur- 
poses of  analysis,  if  the  liability  to  error  in  chemical  manipu- 
lations did  not  make  it  expedient  to  employ  a  variety  of  tests, 
as  corroborative  evidence,  before  coming  to  a  conclusion  in 
regard  to  the  composition  of  a  substance. 

It  is  important  that  the  student  should  turn  to  Part  III.,  and 
commit  to  memory  the  general  features  of  the  scheme  of 
separation  for  each  group,  at  the  time  that  he  is  performing 
the  reactions  of  the  members  of  the  group  as  they  are  described 
in  the  following  pages.  By  this  means  he  will  make  himself 
familiar  with  the  important  points  in  which  the  compounds 
with  which  he  has  to  deal  differ  from  each  other,  and  the 
manner  in  which  these  differences  can  be  used  in  analysis  ;  also 
at  this  stage  of  his  progress  it  is  advisable  for  him  to  make 
mixtures  of  compounds  of  several  elements,  and  to  analyze 
them  according  to  the  directions  given  in  Part  III. 

The  student,  keeping  in  view  the  reasons  for  learning  the 
characteristic  reactions  of  the  compounds  of  each  element, 
should  perform  carefully  the  tests  described  in  Part  II.,  supple- 
menting the  description  by  the  closest  observation  of  the  phe- 
nomena as  they  pass  before  his  eyes.  By  practice  of  this  kind 
he  will  soon  acquire  the  skill  in  manipulation  necessary  for 
analytical  work.  ALWAYS,  WHEN  A  REACTION  is  PERFORMED, 

THE  EQUATION  DESCRIBING  IT  SHOULD  BE  WRITTEN.      The  for- 

28 


PART  II.  29 

mulas  of  the  reagents  and  the  compound  operated  upon*  form 
the  first  half  of  the  equation  ;  the  formula  of  the  precipitate, 
which  is  given  in  the  book,  enters  into  the  second  half  of  the 
equation  and  determines  the  formulas  of  its  other  members. 

Thus  it  is  known  that  baric  chloride  and  calcic  sulphate  give 
a  precipitate  of  BARIC  SULPHATE,  BaSO4.  (See  page  34.) 
From  the  formulas  on  the  labels  of  the  bottles  we  can  con- 
struct the  equation  :  BaCl2  +  CaSO4  =  BaSO4  +  X,  and  by 
inspecting  the  equation  we  find  the  unknown  quantity  X  can 
only  be  CaQ2.  The  following  case  is  more  complicated  :  di- 
sodic  hydric  phosphate,  ammonic  hydrate,  and  magnesic  sul- 
phate form  a  precipitate  of  MAGNESIC  AMMONIC  PHOSPHATE, 
MgNH4PO4 ;  or,  putting  the  statement  into  formulas,  Na^ 
HPO4  +  NH4HO  +  MgSO4  =  MgNH4PO4  +  X.  Here  X  = 
Na2  +  H2  4-  O  -f  SO4,  and  the  question  arises  :  How  are  these 
bodies  combined  ?  A  slight  experience  will  teach  that  the  rule 
2d  (page  9)  brings  the  SO4  and  the  Na  together,  and  conse- 
quently the  H2  and  the  O  ;  while  an  inspection  of  the  tables 
of  quantivalence  (pages  2  and  12)  shows  that  SO4  combines 
with  2Na,  and  that  O  combines  with  2H  ;  hence  X  becomes 
Na2SO4  +  H2O. 

The  grouping  together  of  the  elements  or  groups  of  elements 
appearing  in  reactions  is  not  usually  a  difficult  matter,  and  is 
soon  learned  with  practice. 

The  formulas  of  the  compounds  which  are  most  frequently 
used  in  the  laboratory  stand  after  the  names  of  the  metals  and 
acids,  and  can  be  used  in  writing  equations. 

*  These  formulas  should  be  given  in  full  upon  the  labels  of  the  bottles 
containing  the  compounds  used  and  the  reagents. 


TESTS    FOR   METALS. 


GROUP  I. 

SODIUM,  POTASSIUM,  AND  AMMONIUM. 

THERE  is  no  reagent  which  precipitates  all  the  metals  of  this 
group.  The  salts  of  metals  of  Group  I.  have  a  neutral  reac- 
tion when  they  contain  strong  acids  like  chlorhydric,  nitric, 
and  sulphuric  acids.  They  have  an  alkaline  reaction  when 
they  contain  weak  acids  like  sulphydric,  boracic,  and  carbonic 
acids. 

SODIUM. 
NaCl ;  Na2CO8 ;  NaHO. 

Sodium  compounds  can  be  recognized  by  heating  them  in 
the  loop  of  a  piece  of  fine  platinum  wire  in  the  flame  of  a 
lamp.  Sodium  colors  the  flame  yellow,  and  can  be  recognized, 
even  when  mixed  with  much  larger  quantities  of  other  ele- 
ments, which  alone  impart  other  colors  to  the  flame. 

When  a  liquid  is  to  be  tested,  it  may  be  evaporated  and  the 
residue  brought  on  the  platinum  wire,  or  frequently  it  is  suffi- 
cient to  dip  the  wire  in  the  liquid  and  to  bring  it  into  the  flame 
of  the  lamp. 

No  reagent  *  precipitates  sodium  compounds. 

*  Only  the  reagents  spoken  of  in  this  book  are  referred  to. 

30 


POTASSIUM.  31 

POTASSIUM. 

KC1;  K2SO4. 

Potassium  compounds  color  the  flame  of  a  lamp  violet.  A 
small  admixture  of  sodium  obscures  the  color  of  the  flame  of 
potassium,  but  the  sodium  color  disappears,  and  that  character- 
istic of  potassium  can  be  observed  when  the  flame  is  viewed 
through  a  thick  glass  colored  blue  with  cobalt. 

Mix  together  a  sodium  and  a  potassium  salt,  observe  the 
yellow  color  imparted  to  the  flame  by  the  mixture,  showing 
the  presence  of  sodium,  and  then  examine  the  flame  for  the 
potassium  color  through  a  piece  of  cobalt  glass  thick  enough 
to  exclude  the  sodium  flame.  Examine  a  pure  sodium  flame 
with  cobalt  glass  to  ascertain  that  the  glass  does  not  allow  the 
color  of  the  sodium  to  pass  through  it,  or  until  the  blue  color 
of  the  sodium  flame  can  be  easily  distinguished  from  the  violet 
of  the  potassium  flame.* 

flatinic  Chloride  precipitates  concentrated  solutions 
of  potassic  chloride  as  a  DOUBLE  CHLORIDE  OF  PLATINUM  and 
POTASSIUM,  (KC1)2  PtCl4.  No  other  salt  of  potassium  can  be 
used  for  this  test.  It  is  best  to  evaporate  the  solution  to  dry- 
ness  in  a  water-bath,  with  a  large  quantity  of  platinic  chloride, 
and  to  wash  the  residue  several  times  with  alcohol.  The 
double  chloride  is  left  as  a  yellow  crystalline  powder,  which 
gives  the  potassium  flame.  The  double  chloride  is  slightly 
soluble  in  water,  but  insoluble  in  alcohol. 

No  other  reagent  precipitates  potassium  compounds. 

*  If  the  glass  is  sufficiently*  thick  and  intense  in  color,  the  light  from 
the  sodium  flame  is  completely  excluded.  A  thinner  glass  allows  a  part 
of  the  light  to  pass  through,  but  it  then  has  a  blue  color  which  can  be  dis- 
tinguished after  practice  from  the  violet  color  of  the  potassium  flame, 
which  passes  through  the  glass  with  little  of  its  brilliancy  diminished. 


32  PART  H. 

AMMONIUM. 

NH4C1;  NH4HO. 

Ammonium  compounds  do  not  color  the  flame  of  a  lamp. 
jPlatinic  Chloride  precipitates  ammonic  chloride  as  a 

DOUBLE    CHLORIDE    OF    PLATINUM    AND    AMMONIUM    (NH4C1)2 

PtCl4.  The  precipitate  forms  under  the  same  circumstances, 
and  has  the  same  aspect  and  properties  as  that  obtained  with 
potassic  chloride,  but  it  can  be  distinguished  from  the  latter 
by  the  absence  of  color  imparted  to  the  flame  of  a  lamp  when 
it  is  heated  in  it.  It  is  destroyed  at  a  dull  red  heat,  and  a 
residue  of  metallic  platinum  is  left. 

Sodic  Hydrate,  added  in  excess  to  ammonic  com- 
pounds, causes  them  to  give  off  the  smell  of  AMMONIA  GAS, 
NH3,  especially  when  the  solution  is  heated.  AMMONIA  GAS 
colors  moist  turmeric  paper  brown  (a  delicate  test). 

No  other  compound  interferes  with  the  application  of  this 
test. 


BARIUM. 


GROUP   II. 

BARIUM,  STRONTIUM,  CALCIUM,  AND 
MAGNESIUM. 

THE  chlorides  and  nitrates  of  metals  of  Group  II.  are  soluble 
in  water.  The  sulphates  of  calcium  and  magnesium  are  also 
soluble.  The  solutions  have  a  neutral  reaction  with  test- 
paper. 

Ammonic  and  Sodic  Carbonates  precipitate  the 
metals  of  Group  II.  in  neutral  solutions  as  CARBONATES.  The 
carbonate  of  magnesium  is  very  soluble  in  solutions  of  ammo- 
nic  salts,  particularly  in  ammonic  chloride  ;  therefore  no  pre- 
cipitate of  magnesic  carbonate  is  produced  when  these  salts 
are  present  in  considerable  quantity. 

Central  "Phosphates  (as  disodic  hydric  phosphate) 
precipitate  the  metals  of  Group  II.  in  neutral  or  alkaline  solu- 
tions as  PHOSPHATES. 

The  carbonates  and  phosphates  of  metals  of  Group  II.  are 
soluble  in  dilute  acids,  unless  the  acids  themselves  are  capable 
of  precipitating  the  metals. 

Sulphydric  Acid,  Ammonic  Sulphide,  and 
Ammonic  Hydrate  do  not  precipitate  the  metals  of 
Group  II. 

BARIUM. 

BaCl3;  Ba(NO3)2;  Ba(HO)2. 

Sulphuric  Acid  (dilute)  precipitates  baric  compounds, 
as  BARIC  SULPHATE,  BaSO4,  white  powder. 
3 


34 


PART  II. 


Calcic  Sulphate  and  other  soluble  sulphates  give  the 
same  precipitate  with  baric  compounds. 

Baric  sulphate  is  insoluble  in  acids. 

Ammonic  Oxalate  precipitates  baric  compounds  from 
concentrated  neutral  or  alkaline  solutions,  as  BARIC  OXALATE, 
BaC2O4,  white  powder,  soluble  in  acids. 

Barium  compounds,  particularly  when  moistened  with  chlor- 
hydric  acid,  color  the  flame  yellowish  green. 

STRONTIUM. 

SrCl2;  Sr(N03)2. 

Sulphuric  Acid  (dilute)  and  soluble  sulphates  precipi- 
tate strontium  from  its  solutions,  as  STRONTIUM  SULPHATE, 
SrSO4,  white  powder.  The  precipitate  will  not  appear  imme- 
diately unless  the  solution  be  concentrated.  In  very  dilute 
solutions  it  may  not  appear  at  all,  since  strontium  sulphate  is 
slightly  soluble  in  water  (i  part  in  6900). 

Ammonic  Oxalate  precipitates  strontium  compounds 
from  neutral  or  alkaline  solutions,  as  STRONTIUM  OXALATE, 
SrC2O4,  white  powder,  soluble  in  acids. 

Strontium  compounds,  particularly  when  moistened  with 
chlorhydric  acid,  color  the  flame  brilliant  red. 

CALCIUM. 

CaCl*;  CaS04;  Ca(HO)2. 

Sulphuric  Acid  (dilute)  and  soluble  sulphates,  except 
calcic  sulphate,  precipitate  concentrated  solutions  of  calcium 
compounds,  as  CALCIC  SULPHATE,  CaSO4,  white  powder.  Cal- 
cic sulphate  is  soluble  in  a  considerable  quantity  of  water, 
therefore  no  precipitate  is  produced  in  very  dilute  solutions  by 
sulphuric  acid.  It  is  insoluble  in  dilute  alcohol,  therefore 
sulphuric  acid  (dilute),  with  the  addition  of  a  large  quantity  of 
alcohol,  precipitates  calcic  sulphate  from  even  dilute  solutions 
of  calcium  compounds. 


MAGNESIUM.  2$ 

Ammonic  Oxalate  precipitates  calcic  compounds  from 
neutral  or  alkaline  solutions,  as  CALCIC  OXALATE,  CaC2O4, 
white  powder.  The  precipitate  forms  best  after  standing  some 
time  in  a  solution  to  which  ammonic  hydrate  has  been  added 
in  excess.  Calcic  oxalate  does  not  dissolve  in  acetic  acid. 

Calcium  compounds,  particularly  when  moistened  with  chlor- 
hydric  acid,  color  the  flame  yellowish  red. 

MAGNESIUM. 

MgCl2;  MgS04. 

Hydro-Disodic  JPJiosphate  precipitates  magnesic 
compounds,  to  whose  solution  ammonic  hydrate  and  ammonic 
chloride  have  been  added,  as  MAGNESIC  AMMONIC  PHOSPHATE, 
MgNH4PO4,  white  crystalline  powder,  or  white  flakes  if  the  solu- 
tion is  concentrated.  This  precipitate  only  appears  after  the 
lapse  of  some  time  in  very  dilute  solutions.  It  is  then  crys- 
talline. 

Sodic  Hydrate,  in  excess,  precipitates  magnesic  com- 
pounds, as  MAGNESIC  HYDRATE,  Mg(HO)2,  white  powder,  when 
it  is  boiled  with  their  solutions.  In  case  ammonic  chloride  is 
present,  the  boiling  must  be  continued  until  the  odor  of  ammo- 
nia is  no  longer  perceptible. 


Mix  together  the  chlorides  of  barium,  strontium,  calcium 
and  magnesium.  Add  a  small  quantity  of  ammonic  chloride, 
then  ammonic  hydrate,  and  finally  ammonic  carbonate.  The 
barium,  strontium,  and  calcium  will  be  precipitated  as  carbon- 
ates. Filter,  and  to  the  filtrate  add  hydro-disodic  phosphate. 
The  magnesium  will  be  precipitated  as  magnesic  ammonic 
phosphate.  Wash  the  precipitate,  consisting  of  the  carbonates 
of  barium,  strontium,  and  calcium,  with  water,  and  pour  over  it 
a  small  quantity  of  dilute  hydrochloric  acid.  This  will  decom- 
pose the  carbonates,  and  the  chlorides  of  the  three  metals  will 


36  PART  II. 

be  formed,  which  being  soluble  will  pass  through  the  filter  and 
can  be  caught  in  a  test-tube  or  small  beaker.  Dilute  the  solu- 
tion with  considerable  water  and  add  dilute  sulphuric  acid. 
Barium  and  strontium  will  be  precipitated  as  sulphates,  while 
the  calcic  sulphate  will  remain  in  solution.  Filter,  add  am- 
monic  hydrate  to  the  filtrate  until  the  reaction  becomes  alka- 
line, and  then  ammonic  oxalate.  The  calcium  will  be  precipi- 
tated as  oxalate.  Wash  the  sulphates  of  barium  and  strontium 
on  the  filter,  and  then  moisten  a  small  particle  of  the  mass  with 
chlorhydric  acid  and  examine  it  on  platinum  wire  in  the  flame. 
The  yellowish-green  color  of  barium  will  first  appear.  After 
holding  the  wire  in  the  flame  for  a  few  seconds  dip  it  into 
chlorhydric  acid  and  again  bring  it  into  the  flame.  By  repeat- 
ing this  operation  several  times  the  barium  flame  will  disappear 
and  the  crimson  color,  characteristic  of  strontium,  will  be  very 
plainly  visible.* 

*  Processes  similar  to  the  above  are  used  for  the  separation  of  all  the 
metals  from  each  other,  and  care  must  always  be  taken  to  ascertain  whether 
enough  of  a  reagent  has  been  added  to  completely  effect  a  precipitation 
before  the  next  test  is  proceeded  with  in  the  filtrate,  and  when  a  precipi- 
tate is  to  be  examined,  it  must  be  thoroughly  freed  by  washing  from  the 
solution  which  adheres  to  it. 


ALUMINIUM.  37 


GROUP  III. 

ALUMINIUM  AND  CHROMIUM. 

THE  sulphates,  chlorides,  and  nitrates  of  metals  of  Group  III. 
are  soluble  in  water,  and  the  solution  has  an  acid  reaction  with 
test-paper.  Aluminic  and  chromic  alum  solutions  have  a 
neutral  reaction. 

Ammonic  Hydrate,  Carbonate,  and  Sulphide 
precipitate  the  metals  of  Group  III.  as  HYDRATES. 

Neutral  Phosphates  precipitate  the  metals  of  Group 

III.  as  PHOSPHATES. 

Sulphydric  Acid  does  not  precipitate  the  metals  of 
Group  III. 

ALUMINIUM. 

A12(S04)3 ;  A12C16. 

Ammonic  Hydrate  precipitates  alum inic  compounds 
as  ALUMINIC  HYDRATE,  A12(HO)6,  gelatinous,  white  flakes. 

The  precipitate  is  insoluble  in  ammonic  hydrate  and  in  am- 
monic  chloride,  but  dissolves  in  acids  and  in  sodic  hydrate.  It 
forms  best  on  boiling. 

Sodic  Hydrate  precipitates  aluminic  compounds  like 
ammonic  hydrate,  but  an  excess  of  sodic  hydrate  dissolves  the 
precipitate  so  quickly  that  its  formation  easily  escapes  notice. 

No  precipitate  is  formed  when  the  solution  in  sodic  hydrate 
is  boiled. 

Solid  compounds  of  aluminium  (except  silicates),  when  mois- 
tened with  cobaltic  nitrate  solution,  and  heated  with  the  oxid- 
izing blowpipe  flame  (see  page  25)  on  charcoal,  or  on  a  plat- 
inum wire,  take  a  blue  color. 


38  PART  II. 

CHROMIUM. 

Cr2(S04)3;  Cr2Cl6. 

Solutions  of  chromic  oxide  compounds  are  green. 

Atnmotlic  Hydrate  precipitates  chromic  compounds 
as  CHROMIC  HYDRATE,  Cr2(HO)6,  gelatinous,  dirty  green  flakes. 
The  precipitate  is  insoluble  in  ammonic  hydrate  after  boiling, 
and  in  ammonic  chloride,  but  dissolves  in  acids  and  in  sodic 
hydrate. 

Sodic  Hydrate  precipitates  chromic  compounds  like 
ammonic  hydrate,  but  dissolves  them  when  an  excess  of  sodic 
hydrate  is  present.  CHROMIC  HYDRATE  is  precipitated  from 
its  solution  in  sodic  hydrate  when  the  dilute  solution  is  boiled 
for  some  time. 

^Blowpipe  Reactions.  Compounds  of  chromium 
color  the  borax  bead  green. 

If  chromic  hydrate,  or  any  solid  chromic  compound,  is 
mixed  with  equal  parts  of  sodic  carbonate  and  sodic  nitrate, 
and  heated  red  hot  on  the  platinum  foil,  chromate  of  sodium 
is  formed  by  the  oxidation  of  chromic  oxide.  Chromate  of 
sodium  dissolves  in  water  with  a  yellow  color.  The  color  is 
heightened  when  an  acid  is  added,  and  an  acid  chromate  is 
formed  in  the  solution. 

This  reaction  is  a  characteristic  test  for  chromium  compounds. 


Mix  together  solutions  of  chromium  and  aluminium  salts, 
add  sodic  hydrate  until  the  reaction  becomes  very  strongly 
alkaline  (the  precipitate  which  first  forms  will  dissolve),  dilute 
with  a  considerable  quantity  of  water  in  a  small  flask,  and  boil 
for  several  minutes  after  a  dirty  green  precipitate  has  formed. 
Filter  from  the  precipitate.  This  precipitate  contains  all  the 
chromium.  Test  it  according  to  Part  III.  (102).  The  fil- 
trate contains  all  the  aluminium.  Test  it  according  to  Part 
III.  (104). 


ZINC,  MANGANESE,  AND  IRON.  39 

GROUP    IV. 
ZINC,  MANGANESE,  IRON,  NICKEL,  AND  COBALT. 

THE  sulphates,  chlorides,  and  nitrates  of  metals  of  Group 
IV.  are  soluble  in  water.  The  solutions  have  an  acid  reaction 
with  test-paper. 

Ammonic  Sulphide  precipitates  the  metals  of  Group 
IV.  as  sulphides  ;  if  the  solution  is  not  neutral,  it  should  be 
made  so  with  ammonic  hydrate. 

Sodic  Hydrate  and  Ammonic  Hydrate  pre- 
cipitate the  metals  of  Group  IV.  as  HYDRATES.  The  HYDRATE 
OF  ZINC  is  soluble  in  an  excess  of  the  precipitant,  and  the  HY- 
DRATES OF  NICKEL  AND  COBALT  are  soluble  in  ammonic 
hydrate. 

Sodic  Carbonate  precipitates  the  metals  of  Group  IV. 
as  CARBONATES  (ferric  compounds  as  ferric  hydrate). 

Neutral  Phosphates  precipitate  the  metals  of  Group 

IV.  as  PHOSPHATES. 

Sulphydric  Acid  does  not  precipitate  the  metals  of 
Group  IV.  when  they  are  in  an  acid  solution.  (See  Zinc, 
page  40.) 

SECTION  I. 

ZINC,  MANGANESE,  AND  IRON. 
Metals  whose  sulphides  are  soluble  in  cold  dilute  chlorhydric  acid. 

ZINC. 

ZnSO4;    ZnCl2. 

Metallic  zinc  dissolves  readily  in  dilute  chlorhydric  acid 
and  sulphuric  acid,  with  evolution  of  hydrogen. 


40  PART  II. 

The  metal  melts  readily  when  heated  on  charcoal  with  the 
blowpipe,  and  at  a  high  temperature  it  distils,  and  the  vapor 
burns  with  a  bluish-white  flame,  depositing  an  incrustation  on 
the  charcoal  of  OXIDE  OF  ZINC,  ZnO,  which  is  white  when  hot 
and  yellow  when  cold.  If  the  incrustation  is  moistened  with 
cobaltic  nitrate  and  heated  in  the  oxidizing  flame,  it  turns 
dirty  green.  By  this  test  zinc  can  often  be  recognized  in 
alloys. 

Ammonic  SulpJiide  precipitates  zinc  compounds  as 
the  SULPHIDE  OF  ZINC,  ZnS,  white,  flocculent  precipitate.  The 
sulphide  of  zinc  is  soluble  in  dilute  chlorhydric  acid,  but 
not  in  acetic  acid.  It  is  the  only  white  insoluble  sulphide. 

SulpJiydric  Acid  precipitates  zinc  as  the  SULPHIDE  OF 
ZINC  only  when  the  metal  is  combined  with  acetic  acid.  To 
obtain  the  precipitate,  if  a  stronger  acid  is  present,  add  sodic 
hydrate  until  the  solution  has  a  strongly  alkaline  reaction,  and 
then  add  acetic  acid,  until  the  reaction  becomes  acid,  before 
treating  with  sulphydric  acid. 

Sodic  and  Ammonic  Hydrates  precipitate  zinc 
compounds  as  the  HYDRATE  OF  ZINC,  Zn(HO)2,  white  flakes. 
The  precipitate  dissolves  easily  in  an  excess  of  the  precipi- 
tants  ;  and  the  solution  in  sodic  hydrate  is  not  reprecipitated 
when  it  is  boiled. 

MANGANESE. 

MnSO4 ;  MnCl2. 

Solutions  of  manganese  compounds  have  a  faint  pink  color. 

Ammonic  SulpTlide  precipitates  manganese  com- 
pounds as  the  SULPHIDE  OF  MANGANESE,  MnS,  flesh-colored 
flakes.  The  sulphide  of  manganese  is  soluble  in  dilute  acids. 

Sodic  and  Ammonic  Hydrates  precipitate  man- 
ganese compounds  as  MANGANOUS  HYDRATE,  Mn(HO)2,  white 
flakes,  which  turn  brown  on  exposure  to  the  air.  Manganous 


IRON.  41 

hydrate  is  insoluble  in  an  excess  of  the  precipitant,  but  it  is 
soluble  in  a  large  quantity  of  ammonic  chloride. 

^Blowpipe  Reactions.  Manganese  compounds  color 
the  borax  bead  amethyst  in  the  oxidizing  flame.  If  a  com- 
pound of  manganese  is  heated  with  a  soda  bead  (which  can  be 
made  in  the  same  way  as  a  borax  bead  in  the  loop  of  a  plati- 
num wire)  in  the  oxidizing  flame,  it  colors  it  green,  in  conse- 
quence of  the  formation  of  manganate  of  sodium.  The  same 
color  is  produced  when  a  compound  of  manganese  is  heated 
on  the  platinum  foil  with  sodic  carbonate  and  nitrate.  On 
boiling  the  green  salt  with  water  containing  a  little  alcohol  it 
is  destroyed,  the  color  disappears,  and  brown  flakes  of  man- 
ganic hydrate  are  precipitated. 

IRON  (ferrous  salts). 
FeS04;  FeCl2. 

Metallic  iron  dissolves  readily  in  dilute  acids,  with  evolution 
of  hydrogen. 

Ferrous  salts  in  solution  have  a  pale  green  color. 

Oxidizing  Agents  (as  nitric  acid  and  potassic  chlo- 
rate), when  heated  with  acid  solutions  of  ferrous  salts,  oxidize 
them  to  FERRIC  SALTS,  whose  color  is  brownish  red,  or  reddish 
yellow,  and  is  more  intense  than  the  green  color  of  ferrous  salts. 

Ammonic  Sulphide  precipitates  ferrous  salts  as  FER- 
ROUS SULPHIDE,  FeS,  black  flakes.  Ferrous  sulphide  is  soluble 
in  dilute  acids. 

Sodic  and  Ammonic  Hydrates  precipitate  ferrous 
compounds  as  FERROUS  HYDRATE,  Fe(HO)2,  at  first  nearly 
white,  then  bluish  green,  and,  finally,  by  absorption  of  oxygen,  red- 
dish brown.  Ferrous  hydrate  is  insoluble  in  an  excess  of  sodic 
hydrate.  The  presence  of  a  large  quantity  of  ammonic  salts 
in  a  solution  prevents  its  precipitation. 

Potassic  Ferrocyanide  precipitates  ferrous  com- 
pounds as  POTASSIC  FERROUS  FERROCYANIDE,  K2Fe(FeCv6),* 

*  Cy,  cyanogen,  is  used  as  a  symbol  for  the  group  CN. 


42  PART  II. 

bluish  white ',  turning  quickly  dark  blue,  through  absorption  of 
oxygen  from  the  air. 

Potassic  Ferricyanide  precipitates  ferrous  com- 
pounds, as  TURNBULL'S  BLUE,  Fe8(Fe2Cyi2),*  deep  blue.  This  is 
the  best  test  for  ferrous  compounds. 

The  last  two  precipitates  are  insoluble  in  dilute  acids. 

JPotassic  Sulphocyanate  gives  no  coloration  with 
ferrous  compounds. 

IRON  (ferric  salts). 
Fe2Cl6. 

Reducing  Agents,  as  sulphurous  and  sulphydric  acids, 
metallic  zinc  and  iron,  reduce  ferric  salts  in  solution  to  ferrous 
salts  when  a  free  acid  is  present. 

The  reaction  with  sulphydric  acid  is  accompanied  by  a  pre- 
cipitation of  sulphur,  Fe2Cl6  +  H2S  =  2FeCl2  +  2HC1  +  S. 

A  similar  reaction  takes  place  with  ammonic  sulphide  and 
the  other  salts  of  sulphydric  acid. 

Ferric  salts  in  solution  have  &  yellow  color,  and  they  possess 
a  much  stronger  coloring  power  than  ferrous  salts. 

Ammonic  Sulphide  reduces  ferric  salts  to  ferrous 
salts,  and  then  precipitates  FERROUS  SULPHIDE,  FeS,  black 
flakes. 

Sodic  and  Ammonic  Hydrates  precipitate  ferric 
compounds  as  FERRIC  HYDRATE,  Fe2(HO)6,  red  gelatinous 
flakes,  insoluble  in  an  excess  of  the  precipitants  and  in  am- 
monic salts. 

Potassic  Ferrocyanide  produces  a  precipitate  of 
PRUSSIAN  BLUE,  Fe4(FeCv6)3,  deep  blue,  in  a  solution  of  ferric 
salts. 

JPotassic  Ferricyanide  colors  solutions  of  ferric  salts 
deep  reddish  brown,  but  produces  no  precipitate.  On  the  ad- 
dition of  a  reducing  agent  a  deep  blue  precipitate  forms. 

.  *  See  note  on  preceding  page. 


NICKEL  AND   COBALT.  43 

"Potassic  Sulphocyanate  gives  a  deep  red  color*  with 
the  smallest  traces  of  ferric  compounds  in  acid  solutions. 

The  different  behavior  with  the  last  three  reagents  of  ferrous 
and  ferric  salts  serves  to  distinguish  between  them. 

^Blowpipe  Heaction.  Iron  colors  the  borax  bead 
green  in  the  reducing  flame,  and  reddish  yellow  while  hot,  and 
yellow  while  cold,  in  the  oxidizing  flame. 


Mix  together  solutions  of  zinc,  manganese,  and  iron  salts  ;  if 
a  ferrous  salt  is  taken,  add  a  little  chlorhydric  acid  and  boil 
the  solution  for  a  few  minutes  with  one  or  two  crystals  of  po- 
tassic  chlorate,  in  order  to  convert  the  ferrous  into  ferric  salt. 
Add  sodic  hydrate  to  the  solution  until  the  reaction  is  strongly 
alkaline,  boil  for  a  few  minutes  and  filter.  All  the  manganese 
and  iron  will  be  precipitated.  Test  the  precipitate  for  iron 
according  to  Part  III.  (101),  and  for  manganese  according 
to  Part  III.  (102).  All  the  zinc  will  be  contained  in  the  fil- 
trate ;  test  it  for  zinc  according  to  Part  III.  (103). 


SECTION  II. 

NICKEL  AND  COBALT. 

Metals  whose  sulphides  are  insoluble  in  cold,  dilute  chlorhydric  acid. 

NICKEL. 
NiS04 ;  Ni(N03)2 ;  NiCl* 

Solutions  of  nickel  salts  are  green. 

Ammonic  Sulphide  precipitates  nickel  compounds  as 
the  SULPHIDE  OF  NICKEL,  NiS,  black  flakes.  The  sulphide  of 
nickel  is  insoluble  in  cold,  dilute  chlorhydric  acid.  It  dis- 
solves readily  on  boiling,  or  in  a  strong  acid. 

*  Potassic  sulphocyanate  gives  the  same  color  with  a  solution  containing 
a  large  quantity.of  free  nitric  acid. 


44  PART  II. 

Sodic  and  Ammonic  Hydrates  precipitate  nickel 
compounds  as  the  HYDRATE  OF  NICKEL,  Ni(HO)2,  apple  green. 
The  hydrate  of  nickel  is  insoluble  in  an  excess  of  sodic  hy- 
drate. It  dissolves  in  ammonic  hydrate,  and  the  solution  has 
a  blue  color. 

^Blowpipe  Reactions*  Nickel  colors  the  borax  bead 
in  the  oxidizing  flame  violet,  when  it  is  hot,  and  a  faint  reddish- 
brown  when  it  is  cold.  By  long-continued  reduction  in  the 
reducing  flame,  or  on  charcoal,  the  bead  may  be  obtained  col- 
orless, but  with  gray  specks  of  reduced  metal  in  it. 

COBALT. 

Co(N03)2 ;  CoCl2. 

Solutions  of  cobalt  salts,  when  dilute,  are  red. 

Ammonic  Sulphide  precipitates  cobalt  compounds  as 
the  SULPHIDE  OF  COBALT,  CoS,  black  flakes.  The  sulphide  of 
cobalt  is  insoluble  in  cold  dilute  chlorhydric  acid.  It  dissolves 
readily  on  boiling  or  in  a  strong  acid. 

Sodic  and  Ammonic  Hydrates  precipitate  cobalt 
compounds  at  first  as  a  blue  basic  salt,  which  changes  to  the 
pale  red  COBALTOUS  HYDRATE,  Co(HO)2  on  boiling.  On  ex- 
posure to  the  air  it  becomes  brown,  through  the  formation  of 
cobaltic  hydrate.  Cobaltous  hydrate  is  insoluble  in  an  excess 
of  sodic  hydrate,  but  it  dissolves  in  ammonic  hydrate,  and  the 
solution  is  red,  tinged  with  brown. 

Blowpipe  Reactions.  Cobalt  compounds  color  the 
borax  bead  blue,  and  the  color  is  so  intense  that  a  small  quan- 
tity of  cobalt  eclipses  the  color  produced  by  a  much  larger 
quantity  of  nickel,  when  the  latter  is  mixed  with  it.  The  blue 
color  does  not  disappear  on  reduction,  so  that  when  sufficient 
nickel  is  present  to  hide  the  color  of  a  small  quantity  of  cobalt 
in  a  bead,  the  color  of  the  nickel  may  be  made  to  disap- 
pear by  a  thorough  reduction,  either  on  the  platinum  wire  or 
on  charcoal,  so  that  the  blue  color  characteristic  of  cobalt  can 


COBALT.  45 

be  detected  in  the  bead.  If  the  bead  was  reduced  on  char- 
coal, it  is  advisable  to  remove  it  from  the  charcoal  and  to  melt 
it  on  the  platinum  wire  in  order  to  examine  its  color. 


Mix  together  sulphates  of  zinc,  manganese,  and  nickel,  and 
the  nitrate  of  cobalt  and  ferrous  sulphate,  add  ammonic  hy- 
drate until  a  permanent  precipitate  begins  to  form,  and  then 
ammonic  sulphide  *  until  the  metals  are  completely  precipi- 
tated as  sulphides.  Wash  the  precipitate  on  a  filter  and  treat 
it  with  cold,  dilute  chlorhydric  acid,  in  order  to  separate  the 
sulphides  of  nickel  and  cobalt  from  the  other  sulphides.  See 
Part  III.  (96). 

Test  for  nickel  and  cobalt  according  to  Part  III.  (97)  and 
(OS). 

Test  for  the  other  metals  in  the  chlorhydric  acid  solution 
according  to  Part  III.  (99),  (WO),  (101),  (102),  and 
(103). 

*  In  order  to  precipitate  the  nickel  completely,  the  solution  must  not 
contain  free  ammonic  hydrate,  and  the  ammonic  sulphide  must  not  con- 
tain an  excess  of  ammonic  hydrate.  If  these  precautions  are  not  observed 
a  part  of  the  sulphide  of  nickel  dissolves,  imparting  a  brown  color  to  the 
solution, 


46  PART  II. 


GROUP    F. 

SILVER,   MERCURY,   LEAD,  COPPER,   BISMUTH, 
AND   CADMIUM. 

THE  salts  of  metals  of  Group  V.  with  chlorhydric,  nitric,  and 
sulphuric  acids,  which  are  soluble  in  water,  have  an  acid  re- 
action with  test-paper. 

Sulphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate metals  of  Group  V.  in  neutral  or  acid  solutions  as 
SULPHIDES.  The  sulphides  of  metals  of  Group  V.  are  insolu- 
ble in  AMMONIC  SULPHIDE  and  in  SODIC  HYDRATE,  and  in  dilute 
acids,  even  when  heated  ;  but,  with  the  exception  of  mercuric 
sulphide,  they  are  all  dissolved  when  boiled  with  strong  NITRIC 
ACID. 

Sodic  and  Ammonic  Hydrates  precipitate  metals 
of  Group  V.  as  OXIDES  or  HYDRATES.  The  hydrate  of  lead  is 
somewhat  soluble  in  an  excess  of  sodic  hydrate,  and  the  oxide 
of  silver  and  the  hydrates  of  copper  and  cadmium  are  readily 
soluble  in  an  excess  of  ammonic  hydrate. 

Sodic  Carbonate  precipitates  the  metals  of  Group  V. 

as  CARBONATES. 

Neutral  Phosphates  precipitate  metals  of  Group  V. 
in  neutral  solutions  as  PHOSPHATES, 


SECTION  7. 
SILVER,  MERCUROUS  COMPOUNDS,  AND  LEAD. 

Metals  whose  compounds  are  precipitated  as  chlorides  by  chlor- 
hydric acid. 


MERCUROUS  COMPOUNDS.  47 

SILVER. 
AgN03. 

Metallic  silver  and  its  alloy  with  copper  dissolve  readily  in 
nitric  acid. 
Sulphydric  Acid   and  Ammonic   Sulphide 

precipitate  compounds  of  silver  as  the  SULPHIDE  OF  SILVER, 
Ag2S,  black  flakes. 

Chlorhydric  Acid  precipitates  compounds  of  silver  as 
the  CHLORIDE  OF  SILVER,  AgCl,  white  flakes,  which  settle  readily 
after  boiling  or  prolonged  shaking.  The  chloride  of  silver  dis- 
solves readily  in  ammonic  hydrate.  It  is  insoluble  in  concen- 
trated nitric  acid,  even  when  it  is  boiled  with  it. 

Sodic  and  Ammonic  Hydrates  precipitate  com- 
pounds of  silver  as  the  OXIDE  OF  SILVER,  Ag2O,  a  grayish 
brown  powder. 

The  oxide  of  silver  dissolves  very  readily  in  ammonic  hy- 
drate. 

Blowpipe  Reactions.  Compounds  of  silver,  mixed 
with  sodic  carbonate,  are  easily  reduced  on  charcoal,  by  heat- 
ing in  the  blowpipe  flame,  and  a  hard  globule  of  metallic  silver 
is  obtained. 

MERCUROUS    COMPOUNDS. 

Hg2(N03)2 ;  Hg2Cl2. 

Metallic  mercury  dissolves  readily  in  nitric  acid  'to  form 
mercurous  nitrate. 

Copper  in  the  form  of  a  thin  sheet  or  wire  can  be  coated 
with  mercury  by  immersing  it  in  a  solution  of  mercury  con- 
taining a  free  acid.  The  coating  is  deposited  in  a  longer  or 
shorter  time,  according  to  the  strength  of  the  solution.  It 
takes  the  color  of  mercury  when  it  is  rubbed  gently  with  a  bit 
of  cloth  or  paper.  The  copper  should  be  cleansed  by  immers- 
ing it  in  dilute  nitric  acid  solution  before  it  is  put  in  the  solu- 
tion containing  mercury. 


48  PART  II. 

Oxidi&ing  Agents,  as  chlorine,  aqua  regia,  or  strong 
nitric  acid,  transform  mercurous  into  mercuric  compounds. 
Sulphydric  Acid   and   Ammonic  Sulphide 

precipitate  mercurous  compounds  as  MERCUROUS  SULPHIDE, 
Hg2S,  black.  Mercurous  sulphide  is  not  dissolved  when  it  is 
boiled  with  moderately  strong  nitric  acid.  It  dissolves  readily 
in  aqua  regia. 

Chlorhydric  Acid  precipitates  mercurous  compounds 
as  MERCUROUS  CHLORIDE,  Hg2Cl2,  white  powder.  Mercurous 
chloride  turns  black,  but  does  not  dissolve  when  ammonic  hy- 
drate is  added  to  it.  It  is  insoluble  in  dilute  acids.  It  dis- 
solves in  aqua  regia. 

Sodic  and  Ammonic  Hydrates  give  with  mer- 
curous compounds  black  precipitates,  insoluble  in  an  excess  of 
the  precipitant.  (For  tests  by  heating  mercurous  compounds, 
see  page  50.) 

LEAD. 

Pb(C2H802)2 ;  Pb(N03)2. 

Metallic  lead  dissolves  readily  in  nitric  acid. 

Sulphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate lead  compounds  as  the  SULPHIDE  OF  LEAD,  PbS, 
black.  Sulphide  of  lead  is  oxidized  by  strong  nitric  acid,  with 
formation  of  sulphate  of  lead,  white  powder y  which  is  insoluble, 
unless  a  very  large  quantity  of  nitric  acid  is  present. 

Chlorhydric  Acid  precipitates  lead  compounds  as 
the  CHLORIDE  OF  LEAD,  PbCl2,  white.  When  the  solution  is 
dilute  no  precipitate  is  produced.  The  chloride  of  lead  dis- 
solves entirely  on  boiling  with  a  large  quantity  of  water. 

Sulphuric  Acid  precipitates  lead  compounds  as  the 
SULPHATE  OF  LEAD,  PbSO*,  white  powder.  Sulphate  of  lead 
is  soluble  to  some  extent  in  chlorhydric  and  nitric  acids,  and 
it  is  slightly  soluble  in  water.  It  is  insoluble  in  a  mixture  of 
alcohol  and  water.  When  complete  precipitation  is  required, 
it  is  best  to  evaporate  with  an  excess  of  sulphuric  acid  until 


LEAD.  49 

all  the  other  acids  are  driven  off,  then  to  dilute  with  water, 
and  to  add  an  equal  bulk  of  alcohol. 

LEAD,  BARIUM,  STRONTIUM,  CALCIUM,  and  STANNIC  COM- 
POUNDS are  the  only  ones  precipitated  by  SULPHURIC  ACID. 

Sodic  and  Ammonic  Hydrates  precipitate  com- 
pounds of  lead  as  BASIC  SALTS  OF  LEAD,  white  precipitate. 
Basic  salts  of  lead  are  somewhat  soluble  in  sodic  hydrate. 

blowpipe  Reactions.  Solid  compounds  of  lead  give, 
when  heated  with  sodic  carbonate  on  charcoal,  globules  of 
metallic  lead,  recognizable  by  their  softness  and  ductility.  An 
incrustation  of  OXIDE  OF  LEAD,  PbO,  deep  yellow  when  hot,  light 
yellow  when  cold,  is  formed  upon  the  charcoal,  not  far  from  the 
place  where  the  substance  is  heated. 

Pure  lead,  or  alloys  containing  a  large  amount  of  lead,  when 
heated  on  charcoal,  without  soda,  burn  with  a  blue  flame,  and 
give  the  PbO  incrustation. 


Mix  together  acetate  of  lead,  nitrate  of  silver,  and  mercurous 
nitrate,  add  dilute  chlorhydric  acid  until  a  precipitate  ceases 
to  form  on  further  addition,  and  filter  the  liquid.  The  pre- 
cipitate contains  all  the  silver  and  mercury  and  a  part  of  the 
lead  as  chlorides.  Add  to  the  filtrate  an  equal  bulk  of  alcohol 
and  a  small  quantity  of  dilute  sulphuric  acid.  All  the  lead 
which  it  contains  will  be  precipitated.  Filter,  wash  the  pre- 
cipitate, and  test  it  for  lead  on  charcoal. 

Make  a  hole  in  the  bottom  of  the  filter  and  wash  the  pre- 
cipitate obtained  with  chlorhydric  acid  into  a  small  flask. 
Test  the  precipitate  for  the  remainder  of  the  lead,  and  for 
silver  and  mercury,  according  to  Part  III.  (#7),  (08),  and 
(69). 

4 


PART  II. 


SECTION  II. 

MERCURIC  COMPOUNDS,  COPPER,  BISMUTH,  AND  CADMIUM. 
Metals  which  are  not  precipitated  by  chlorhydric  acid. 

MERCURIC    COMPOUNDS. 

HgCl2. 

Snlphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate mercuric  compounds,  at  first  as  double  salts,  which 
appear  first  white,  then  yellow,  orange,  and  brown,  and  finally  as 
MERCURIC  SULPHIDE,  HgS,  black.  Mercuric  sulphide  does  not 
dissolve  when  it  is  boiled  with  moderately  concentrated  nitric 
or  chlorhydric  acid.  It  dissolves  readily  in  aqua  regia. 

Sodic  Hydrate  precipitates  mercuric  compounds  at 
first  as  basic  salts,  reddish  brown,  finally  as  MERCURIC  OXIDE, 
HgO,  yellow  ;  insoluble  in  an  excess  of  the  precipitant. 

Ammonic  Hydrate  precipitates  mercuric  compounds 
as  SALTS  CONTAINING  AMMONIA,  white.  The  precipitate  is 
insoluble  in  an  excess  of  ammonic  hydrate. 

StannoMS  Chloride  reduces  mercuric  compounds,  and 
precipitates  them  as  MERCUROUS  CHLORIDE,  Hg2Cl2,  white. 
After  the  metals  of  Group  V.,  Section  I.,  if  they  are  present 
have  been  removed  from  a  solution  by  the  addition  of  chlorhy- 
dric acid,  mercuric  compounds  are  the  only  ones  which  give  a 
precipitate  in  acid  solution  with  stannous  chloride. 

Reactions  with  the  Aid  of  Heat.  Dry  mercurous 
and  mercuric  chlorides  form  white  sublimates  when  they  are 
heated  in  a  closed  glass  tube.  All  dry  compounds  of  mer- 
cury, when  they  are  mixed  with  dry  sodic  carbonate,  and 
heated  in  a  closed  tube,  give  a  sublimate  of  metallic  mercury. 
The  sublimate  is  at  first  a  faint  metallic  film,  which  augments 
until  drops  of  mercury  appear.  If  the  quantity  of  mercury  is 


COPPER. 


small,  the  film  may  be  made  to  take  the  form  of  metallic  drops 
by  rubbing  it  with  a  copper  wire. 


COPPER. 
CuSCV,  CuCl2. 

Metallic  copper  dissolves  readily  in  dilute  nitric  acid.  It 
dissolves  with  difficulty  in  chlorhydric  acid.  All  copper  solu- 
tions are  blue  or  green. 

Iron  OT  Zinc  precipitates  copper  from  its  acid  solutions, 
either  as  a  metallic  coating  or  as  brownish  red  metallic  grains. 
If  a  strip'of  zinc  and  one  of  platinum  are  placed  in  a  dilute 
acid  solution  of  copper,  so  that  they  touch  each  other,  the 
platinum  is  plated  with  copper. 

Sulphydric  Acid  and  Ammonic  Sulphide 
precipitate  cupric  compounds  as  CUPRIC  SULPHIDE,  CuS,  black. 
Cupric  sulphide  is  insoluble  in  dilute  acids,  but  dissolves  readily 
in  strong  acids.  It  is  somewhat  soluble  in  an  excess  of  am- 
monic  sulphide. 

Sodic  Hydrate  precipitates  cupric  compounds  as 
CUPRIC  HYDRATE,  Cu(HO)2,  light  blue.  On  boiling,  CUPRIC 
OXIDE,  CuO,  black,  is  formed.  The  precipitate  is  insoluble  in 
an  excess  of  sodic  hydrate. 

Ammonic  Hydrate  precipitates  cupric  compounds  as 
CUPRIC  HYDRATE,  which  dissolves  immediately  in  an  excess  of 
ammonic  hydrate.  The  solution  has  a  very  intense  blue  color. 

A  valuable  test  for  cupric  compounds. 

Ferrocyanide  of  Potassium  precipitates  cupric 
compounds  in  acid  solutions,  as  FERROCYANIDE  OF  COPPER, 
Cu2(FeCy6),  reddish  brown.  This  is  a  delicate  test  for  very 
small  quantities  of  copper. 

Blowpipe  Reactions*  Cupric  compounds  color  the 
borax  bead  blue  when  cold,  and  green  when  hot,  in  the  oxidiz- 
ing flame.  The  bead  is  colored  red,  and  becomes  opaque  in 


ij2  PART  II. 

the  reducing  flame.  No  other  metal  produces  this  color- 
ation. 

When  cupric  compounds  are  heated  with  sodic  carbonate 
on  charcoal,  metallic  copper  is  reduced  in  the  form  of  small 
globules,  which  can  be  easily  recognized  by  their  hardness  and 
red  color. 

Copper  Flame.  Compounds  containing  copper  (alloys 
and  salts)  color  the  flame  of  a  lamp  green,  or  if  chlorine  is 
present,  blue.  By  moistening  a  cupric  compound  with  chlorhy- 
dric  acid  the  blue  color  can  easily  be  detected. 

BISMUTH. 
Bi(N03)3. 

Metallic  bismuth  dissolves  readily  in  moderately  concen- 
trated nitric  acid.  It  dissolves  with  great  difficulty  in  chlor- 
hydric  acid. 

Water,  Solutions  of  bismuth,  particularly  those  contain- 
ing chlorhydric  acid,  are  remarkable  for  giving  a  precipitate 
consisting  of  a  BASIC  SALT,  when  water  is  added  to  them, 
unless  they  contain  a  large  quantity  o£  free  acid.  The  basic  salt 
can  be  dissolved  by  the  addition  of  an  acid.  CHLORHYDRIC 
ACID  precipitates  nitrate  of  bismuth  solution  as  a  BASIC  CHLOR- 
IDE because  the  basic  chloride  of  bismuth  is  more  insoluble 
than  the  other  basic  salts.  .  The  precipitate  is  soluble  on  the 
further  addition  of  chlorhydric  acid. 

Sulphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate bismuth  compounds  as  the  SULPHIDE  OF  BISMUTH, 
Bi3S8,  black. 

Sodic  and  Ammonic  Hydrates  precipitate  bis- 
muth compounds  as  the  HYDRATE  OF  BISMUTH,  Bi(HO)3,  white. 

The  hydrate  of  bismuth  is  insoluble  in  an  excess  of  sodic 
and  ammonic  hydrates. 

Blowpipe  Reactions.  Bismuth  compounds,  mixed 
with  sodic  carbonate,  and  heated  on  charcoal,  give  brittle 


CADMIUM. 


53 


metallic  globules,  and  an  incrustation  of  OXIDE  OF  BISMUTH, 
Bi2O3,  on  the  charcoal. 

The  incrustation  is  orange  yellow  when  hot,  and  bright  yellow 
when  cold. 

CADMIUM. 

Cd. 

Metallic  cadmium  dissolves  readily  in  nitric  acid,  and  the 
solution  contains  nitrate  of  cadmium,  Cd(NO8)2. 
Sulphydric  Acid   and  Ammonic  Sulphide 

precipitate  cadmium  compounds  as  the  sulphide,  CdS,  lemon 
yellow.  This  is  insoluble  in  ammonic  sulphide  and  potassium 
cyanide,  but  readily  soluble  in  hot  dilute  sulphuric  or  nitric 
acid. 

Sodic  and  Ammonic  Hydrates  precipitate  com- 
pounds of  cadmium  as  the  hydrate,  Cd(HO)2,  very  readily  solu- 
ble in  a  slight  excess  of  ammonic  hydrate,  but  insoluble  in 
sodic  hydrate. 

When  compounds  of  cadmium  are  mixed  with  sodic  carbon- 
ate and  exposed  on  charcoal  to  the  inner  blowpipe  flame,  the 
charcoal  becomes  covered  with  an  incrustation  of  yellow  or 
brownish-yellow  oxide  of  cadmium,  CdO. 


Mix  together  mercuric  chloride,  cupric  sulphate,  and  the 
nitrates  of  bismuth  and  cadmium,  and  add  sulphydric  acid 
until  the  metals  are  completely  precipitated.  Filter,  and  wash 
the  precipitates  with  a  little  water ;  spread  the  filter  out ; 
scrape  the  precipitate  from  it,  and  heat  the  precipitate  gently 
in  a  porcelain  dish  with  strong  nitric  acid  until  red  fumes 
cease  to  be  given  off ;  then  add  a  little  water  and  boil  for  a  few 
minutes.  The  sulphide  of  mercury  remains  insoluble.  Filter 
it  off  and  test  according  to  Part  III.  (89).  Test  the  filtrate 
for  bismuth,  copper,  and  cadmium  according  to  Part  III.  (91), 
(92),  and  (93). 


54  PART  II. 


GROUP    VI. 
TIN,  ANTIMONY,  ARSENIC,  AND  GOLD. 

THE  metals  of  Group  VI.  sometimes  act  as  acids,  uniting 
with  metals,  and  sometimes  as  metals,  uniting  with  acids. 
Their  combinations  with  acids  have  an  acid  reaction  ;  those 
with  metals  which  are  soluble  in  water  have  an  alkaline  re- 
action. 

Sulphydric  Acid  precipitates  the  metals  of  Group  VI. 
as  SULPHIDES.  The  precipitation  takes  place  slowly,  particu- 
larly in  the  case  of  arsenic  acid,  and  it  should  never  be  con- 
sidered complete  unless  a  current  of  sulphydric  acid  gas  is 
passed  through  the  solution  for  some  time,  and  it  is  left  to 
stand  twenty-four  hours. 

The  sulphides  of  metals  of  Group  VI.  are  insoluble  in  dilute 
acids,  but  they  are  decomposed  or  dissolved  by  boiling  with 
concentrated  acids.  They  are  soluble  in  sodic  hydrate  and 
in  ammonic  sulphide  (the  sulphide  of  gold  with  difficulty). 

TIN. 

Sn. 

Metallic  tin  dissolves  readily  in  strong  chlorhydric  acid  on 
boiling,  and  the  solution  contains  STANNOUS  CHLORIDE,  SnCl3. 
Tin  is  oxidized  to  STANNIC  OXIDE,  SnO2,  white  powder,  by 
strong  nitric  acid.  Stannic  oxide  is  insoluble  in  nitric  acid, 
but  dissolves  readily  in  hot  concentrated  chlorhydric  acid, 
and  the  solution  contains  STANNIC  CHLORIDE,  SnCl4.  Stannic 
chloride  is  also  formed  by  the  action  of  aqua  regia  on  metallic 
tin. 


STANNOUS  COMPOUNDS. 


55 


Zinc  precipitates  METALLIC  TIN  from  its  solu- 
tions in  acids  as  crystalline  metallic  particles. 

^Blowpipe  ^Reactions.  Compounds  of  tin  are  re- 
duced, when  they  are  heated  with  sodic  carbonate  and  potas- 
sic  cyanide  on  charcoal  ;  and  METALLIC  TIN  may  be  discov- 
ered in  flattened  globules  by  rubbing  the  fused  mass,  taken 
from  the  charcoal,  in  a  mortar  with  water,  and  washing  several 
times  by  decantation.  When  the  globules  are  large,  they  may 
be  observed  on  the  charcoal  during  the  fusion. 

Oxidation  by  Sodic  Nitrate.  When  a  sulphide  of 
tin  is  oxidized  at  the  lowest  possible  temperature  by  a  mixture 
of  sodic  nitrate  and  carbonate,  STANNIC  OXIDE,  SnO2,  which 
is  insoluble  in  water,  is  formed.  If  the  oxidation  is  carried  on 
at  too  high  a  temperature,  stannate  of  sodium  is  formed,  which 
is  soluble  in  water. 


STANNOUS   COMPOUNDS. 

SnCl2. 

Snlphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate stannous  compounds  from  acid  solutions  as  STANNOUS 
SULPHIDE,  SnS,  dark  brawn.  Stannous  sulphide  dissolves  with 
difficulty  in  a  colorless  solution  of  ammonic  sulphide  (the 
mono- sulphide),  and  it  is  scarcely  soluble  in  ammonic  hydrate 
and  carbonate.  It  is  converted  into  stannic  sulphide  by  a 
yellow  solution  of  ammonic  sulphide  (the  poly-sulphide),  and 
dissolves  readily  when  warmed  with  the  solution.  Stannous 
sulphide  is  soluble  in  sodic  hydrate.  The  sulphides  of  tin  are 
precipitated  from  these  solutions,  when  a  dilute  acid  is  added 
gradually,  until  the  reaction  becomes  strongly  acid. 

Sodic  and  Ammonic  Hydrates  and  Carbo- 
nates precipitate  stannous  compounds  as  STANNOUS  HY- 
DRATE, Sn(HO)2,  white. 

Stannous  hydrate  dissolves  in  an  excess  of  sodic  hydrate, 


56  PART  77. 

but  it  is  insoluble  in  an  excess  of  sodic  carbonate  and  of  am- 
monic  hydrate  and  carbonate. 

Mercuric  Chloride  changes  stannous  chloride,  SnCl2, 
into  STANNIC  CHLORIDE,  SnCl4,  and  a  white  precipitate  of  MER- 
CUROUS  CHLORIDE,  Hg2Cl2,  is  formed.  No  other  metal  in  solu- 
tion gives  this  reaction  with  mercuric  compounds. 

STANNIC   COMPOUNDS. 

SnCl<. 

Swlphydric  Acid   and  Ammonic   Sulphide 

precipitate  stannic  compounds  as  STANNIC  SULPHIDE,  SnS2, 
light  yellow.  Stannic  sulphide  dissolves  readily  in  ammonic 
sulphide,  and  in  sodic  hydrate,  and  is  precipitated  completely 
from  the  solution,  when  a  dilute  acid  is  added  gradually  until 
the  reaction  becomes  strongly  acid.  Stannic  sulphide  is  nearly 
insoluble  in  ammonic  carbonate. 

Sodic  and  Ammonic  Hydrates  and  Carbon- 
ates precipitate  stannic  compounds  as  STANNIC  HYDRATE, 
Sn(HO)4,  white.  Easily  soluble  in  sodic  hydrate. 

Mercuric  Chloride  does  not  give  a  precipitate  with 
stannic  chloride. 

Metallic  Zinc  precipitates  TIN  as  crystalline  metallic 
particles  from  stannic  compounds  in  an  acid  solution,  and  the 
tin  can  be  easily  recognized  by  dissolving  the  metallic  parti- 
cles, after  they  have  been  washed  by  decantation.  They  are 
dissolved  by  heating  them  with  a  few  drops  of  strong  chlorhy- 
dric  acid,  and  the  stannous  chloride  thus  obtained  gives  the 
precipitate,  above  described,  with  mercuric  chloride. 

ANTIMONY. 

Sb. 

Metallic  antimony  is  scarcely  attacked  by  chlorhydric  acid. 
It  is  oxidized  when  heated  with  moderately  strong  nitric  acid 


ANTIMONY.  57 

to  ANTIMONIC  OXIDE,  Sb2O5,  which  is  almost  completely  inso- 
luble in  nitric  acid,  but  is  readily  soluble  in  hot  concentrated 
chlorhydric  acid.  Aqua  regia  dissolves  metallic  antimony  as 

ANTIMONIC  CHLORIDE,  SbCl5. 

Water  precipitates  solutions  of  antimony  containing  chlor- 
hydric acid,  particularly  those  of  antimonious  compounds,  as 
a  BASIC  CHLORIDE.  The  precipitation  can  be  prevented,  or 
the  precipitate  dissolved  by  the  addition  of  a  sufficient  quan- 
tity of  acid  ;  for  this  purpose  tartaric  acid  is  the  most  suitable. 

Metallic  Zinc  partially  precipitates  METALLIC  ANTI- 
MONY from  its  acid  solutions,  and  when  a  piece  of  platinum 
in  contact  with  a  piece  of  zinc  is  introduced  into  solutions  of 
antimony  containing  an  excess  of  chlorhydric  acid,  metallic 
antimony  is  deposited  upon  the  platinum  as  a  dark-brown  stain. 
No  other  metal  produces  the  same  stain  under  like  circumstances. 

Antimoniuretted  Hydrogen.  To  obtain  this  body, 
follow  exactly  the  directions  given  for  obtaining  arseniuretted 
hydrogen.  (See  page  59.) 

The  mirror  obtained  with  antimoniuretted  hydrogen  consists 
of  a  black,  sooty,  metallic  deposit.  It  dissolves  very  slowly  in 
hypochlorite  of  sodium.  If  the  deposit  is  moistened  with 
yellow  ammonic  sulphide,  an  orange-yellow  stain  appears  on  the 
spot  when  it  is  dried.  These  reactions  are  unimportant  as 
tests  for  antimony,  but  a  knowledge  of  them  is  necessary,  in 
order  that  they  may  not  be  mistaken  for  evidences  of  the  pres- 
ence of  arsenic. 

^Blowpipe  Reactions*  Antimony  compounds,  mixed 
with  sodic  carbonate  and  potassic  cyanide,  are  quickly  reduced, 
at  a  comparatively  low  temperature,  to  METALLIC  ANTIMONY, 
brittle  shining  grains.  The  metal  is  completely  volatilized  by 
a  strong  heat.  Metallic  antimony,  when  heated  on  charcoal, 
burns  with  a  white  smoke,  and  gives  an  incrustation  which  is 
deposited  at  some  distance  from  the  place  heated,  and  is  very 
volatile  before  the  blowpipe  flame. 

Oxidation  by  Sodic  Nitrate.    When  a  sulphide  of 


5  8  PART  II. 

antimony  is  oxidized  by  a  mixture  of  sodic  carbonate  and 
nitrate,  ANTIMONIATE  OF  SODIUM,  Na3SbO4,  insoluble  in  water, 
is  formed. 

ANTIMONIOUS  COMPOUNDS. 

SbCl3 ;  KSbOC4H4O6,  Tartar  Emetic. 

Sulphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate antimonious  compounds  as  ANTIMONIOUS  SULPHIDE, 
Sb2S8,  orange  red.  The  precipitation  should  be  made  in  a  cold 
solution  containing  tartaric  acid  and  very  little  free  chlorhy- 
dric  acid.  Antimonious  sulphide  dissolves  in  sodic  hydrate 
and  in  ammonic  sulphide,  most  readily  in  yellow  ammonic  sul- 
phide, and  is  precipitated  completely  from  these  solutions  when 
a  dilute  acid  is  added  gradually  until  the  reaction  becomes 
acid.  It  is  nearly  insoluble  in  ammonic  carbonate.  Yellow 
ammonic  sulphide  converts  it  into  ANTIMONIC  SULPHIDE,  Sb2S6. 
(See  below.) 

Sodic  and  Ammonic  Hydrates  and  Carbon- 
ates precipitate  antimonious  compounds  as  ANTIMONIOUS 
OXIDE,  Sb2O3,  voluminous  white  precipitate.  The  precipitation 
only  takes  place  after  the  lapse  of  a  considerable  time  in  solu- 
tions containing  tartaric  acid.  Antimonious  oxide  is  soluble  in 
sodic  hydrate. 

ANTIMONIC   COMPOUNDS. 

SbCl5;  Sb206. 

Sulphydric  Acid  and  Ammonic  Sulphide  pre- 
cipitate antimonic  compounds  in  acid  solution  as  ANTIMONIC 
SULPHIDE,  Sb2S6,  orange  red.  The  precipitation  does  not  take 
place  immediately  ;  but  first  an  orange  color  appears  ;  and  it  is 
only  after  passing  sulphydric  acid  for  a  long  time  through  the 
solution  that  all  the  antimonic  sulphide  is  precipitated.  The 
precipitation  should  be  made  in  a  cold  solution  containing  tar- 


ARSENIC. 


59 


taric  acid  and  very  little  free  chlorhydric  acid.  Antimonic 
sulphide  has  the  same  properties  as  antimonious  sulphide. 

Antimonic  oxide,  Sb2O5,  plays  the  part  of  an  acid  with  bases, 
and  forms  insoluble  compounds  with  SODIUM,  and  soluble 
compounds  with  POTASSIUM. 

Compounds  of  antimony  can  best  be  recognized  by  the  stain, 
which  they  form  on  platinum,  when  the  metal  is  precipitated 
by  zinc  from  their  solutions.  (See  page  57,  Antimony.) 

ARSENIC. 
As. 

Metallic  arsenic  is  readily  oxidized  to  arsenious  or  to  ar- 
senic compounds  by  nitric  acid  and  is  dissolved.  It  is  not  dis- 
solved by  chlorhydric  acid.  Metallic  arsenic  volatilizes  at  a 
heat  below  redness  in  a  tube  or  on  charcoal,  and  produces  an 
odor  like  garlic,  very  characteristic  of  compounds  of  arsenic. 
The  same  odor  is  produced  by  heating  a  dry  compound  of 
arsenic  on  charcoal. 

Arseniuretted  Hydrogen.  To  obtain  this  body, 
and  to  use  Marsh's  test,  provide  a  4  oz.  flask  with  a  tight- 
fitting  cork,  into  which  a  funnel-tube  and  a  small  tube  drawn 
off  to  a  point  and  bent  at  a  right  angle  *  are  introduced.  Fill 
it  two-thirds  full  of  dilute  sulphuric  acid,  and  add  several 
pieces  of  pure  zinc.  A  brisk  evolution  of  hydrogen  only  com- 
mences after  a  few  minutes.  When  that  point  is  reached, 
wait  five  minutes  for  the  expulsion  of  the  air  contained  in  the 
flask  (without  this  precaution  there  is  danger  of  an  explosion), 
and  light  the  hydrogen  issuing  from  the  point  of  the  bent  tube. 
The  opening  of  the  point  and  the  quantity  of  gas  evolved 
should  be  such  that  the  hydrogen  burns  with  a  blunt  flame 

*  In  order  to  dry  the  gas,  it  is  better  to  fit  to  the  cork  a  chloride  of  cal- 
cium tube,  bent  downwards,  so  that  it  will  not  tip  the  flask  over  by  its 
weight,  and  to  adapt  to  the  chloride  of  calcium  tube,  by  means  of  a  cork, 
a  suitable  jet  directed  upwards. 


60  PART  II. 

about  J  in.  in  length.  Hold  a  bit  of  porcelain  in  the  flame,  in 
order  to  be  certain  that  the  reagents  employed  are  free  from 
arsenic.  (See  below.)  If  this  is  the  case,  pour  a  little  of  an 
acid  liquid  containing  arsenic  into  the  funnel-tube,  while  the 
hydrogen  flame  continues  lighted.  The  flame  after  a  few 
moments  becomes  white,  and  leaves  a  black  stain  of  metallic 
arsenic  upon  a  cold  porcelain  object  held  in  it,  in  the  same 
way  that  the  flame  of  a  candle  would  leave  a  deposit  of  soot 
on  a  cold  surface.  The  arsenic  stain  or  mirror  is  shining  black 
(not  sooty  black,  like  antimony).  It  dissolves  quickly  in  hypo- 
chlorite  of  sodium.  When  the  stain  is  moistened  with  yellow 
ammonic  sulphide  solution,  and  the  spot  is  dried,  it  is  bright 
yellow. 

Reaction  with  the  Aid  of  Heat.  All  dry  com- 
pounds of  arsenic,  except  some  compounds  with  the  higher 
metals,  when  they  are  mixed  with  dry  sodic  carbonate  and  po- 
tassic  cyanide,  and  heated  in  a  sealed  tube,  give  a  sublimate 
of  metallic  arsenic,  which  can  be  best  recognized  by  breaking 
the  sealed  end  of  the  tube  after  the  formation  of  the  sublimate, 
and  by  heating  the  sublimate  quickly  until  it  begins  to  volatil- 
ize, and  by  smelling  of  the  upper  end  of  the  tube.  A  smell 
of  garlic  is  proof  of  the  presence  of  arsenic. 


ARSENIOUS  COMPOUNDS. 

As2O8. 

Arsenious  oxide  is  sparingly  soluble  in  water  It  dissolves 
more  readily  in  sodic  hydrate  or  carbonate,  or  in  chlorhydric 
acid. 

Sulphydric  Acid  and  Ammonic  SulpJiide  pre- 
cipitate arsenious  compounds  from  acid  solutions  as  ARSENI- 
OUS SULPHIDE,  As2S8,  yellow.  Arsenious  sulphide  dissolves 
very  readily  in  sodic  hydrate  and  ammonic  sulphide,  and  it 
also  dissolves,  although  less  readily,  in  ammonic  carbonate. 


ARSENIC  COMPOUNDS.  61 

It  is  precipitated  from  these  solutions,  when  a  dilute  acid  is 
slowly  added  until  the  reaction  becomes  acid. 

Sodic  and  Ammonic  Hydrates  and  Carbon- 
ates produce  no  precipitate  in  arsenious  compounds. 

ARSENIC   COMPOUNDS. 

As2O5. 

By  heating  any  compound  of  arsenic  with  strong  nitric  acid, 
or  by  fusing  a  dry  compound  with  a  mixture  of  dry  sodic  car- 
bonate and  nitrate  in  a  porcelain  crucible,  arsenic  acid, 
H3AsO4,  or  arseniate  of  soda,  Na^AsC)^  is  formed.  These  com- 
pounds are  soluble  in  water. 

Sulphydric  Acid  and  Ammonic  Sulphide  do 
not  immediately  precipitate  acid  solutions  of  arsenic  acid. 

The  solution  (which  should  contain  free  chlorhydric  acid) 
at  first  becomes  yellow,  when  a  current  of  sulphydric  acid  is 
passed  through  it,  and  finally  a  yellow  precipitate  is  formed  ; 
reduction  to  the  state  of  an  arsenious  compound  takes  place 
slowly,  and  at  the  end  of  several  days  all  the  arsenic  is  pre- 
cipitated as  ARSENIOUS  SULPHIDE,  As2S3,  yellow.  This  opera- 
tion may  be  very  materially  accelerated  by  boiling  the  solution. 

See  the  properties  of  arsenious  sulphide  above. 

Argentic  ^Nitrate  produces  no  precipitate  in  acid 
solutions  of  arsenic  compounds.  {When  chlorhydric  acid  is 
present  a  white  precipitate  of  argentic  chloride  forms,  and 
must  usually,  after  the  addition  of  an  excess  of  argentic  nitrate, 
be  separated  by  decantation  and  filtration.)  If  sodic  hydrate 
is  added  to  a  clear  solution  containing  arsenic  acid,  and  con- 
taining also  an  excess  of  argentic  nitrate,  until  a  permanent 
precipitate  forms,  and  if  then  acetic  acid  is  added  until  the 
reaction  becomes  acid,  ARGENTIC  ARSENIATE,  Ag3AsO4,  brick 
red,  is  precipitated.  This  precipitate  is  soluble  in  ammonic 
hydrate  and  in  dilute  nitric  acid. 

This  is  the  usual  test  for  arsenic. 


62  PART  II. 

GOLD. 

AuCl8. 

Metallic  gold  is  insoluble  in  any  single  acid,  but  it  dissolves 
readily  in  aqua  regia. 

Ferrous  Sulphate  solution  in  considerable  quantity 
precipitates  METALLIC  GOLD,  Au,  purplish-brown  powder,  from 
acid  solutions  containing  gold. 

Stan/nous  Chloride  solution  precipitates  METALLIC 
GOLD,  Au,  purple  flakes,  very  finely  divided,  from  acid  solutions 
containing  gold.  (A  very  dilute  solution  of  stannous  chloride 
should  be  added,  drop  by  drop,  to  the  gold  solution.) 

Sulphyciric  A.d(l  gas  slowly  precipitates  gold  com- 
pounds in  acid  solutions  as  SULPHIDE  OF  GOLD,  Au2S3,  brown 
flakes.  The  precipitate  is  soluble  after  a  long  digestion  with 
yellow  ammonic  sulphide. 

^Blowpipe  Reaction.  Gold  is  easily  separated  from 
its  compounds  with  non-metallic  elements  by  heating  them  on 
charcoal.  It  can  be  recognized  as  bright  yellow  globules. 

Gold  can  easily  be  separated  from  the  other  metals  and  recog- 
nized by  its  precipitation  with  ferrous  sulphate,  or  by  its  insolu- 
bility in  any  single  acid. 


Mix  together  solutions  containing  tin,  antimony,  and  arsenic, 
dilute  the  solution  and  add  a  considerable  quantity  of  dilute 
chlorhydric  acid.  Pass  sulphydric  acid  through  the  solution 
for  an  hour,  and  set  it  over  night  in  a  warm  place.  Filter  the 
precipitate,  wash  and  dry  it  completely.  It  contains  all  the 
metals  in  the  form  of  sulphides  ;  separate  them  according  to 
Part  III.  (83)  and  the  following  tests. 


TESTS  FOR  ACIDS. 


GROUP    I. 

ARSENIOUS,  ARSENIC,  CHROMIC,  SULPHURIC, 
PHOSPHORIC,  BORACIC,  OXALIC,  FLUORHY- 
DRIC,  CARBONIC,  AND  SILICIC  ACIDS. 

Acids  which  are  precipitated  from  neutral  or  slightly  alkaline 
solutions  by  baric  chloride. 

SECTION  I. 

ARSENIOUS,  ARSENIC,  AND  CHROMIC  ACIDS. 

Acids  which  are  precipitated  as  sulphides  or  reduced  to  an  oxide 
by  sulphydric  acid. 

ARSENIOUS  and  ARSENIC  acids  are  precipitated  as  sulphides, 
and  must  always  be  detected  by  sulphydric  acid. 

Arsenic  with  this  reagent  plays  the  part  of  a  metal.  (See 
pages  60  and  61.) 

CHROMIC    ACID. 

H2CrO4  ;  K2Cr2O7. 

Solutions  containing  chromic  acid  have  a  yellow  color. 
Those  which  have  an  acid  reaction  are  of  a  deeper  color,  and 

63 


64  PART  II. 

very  concentrated  acid  solutions  are  red.  Solutions  contain- 
ing chromic  acid  can  usually  be  recognized  by  these  colors. 
Chromic  acid  can  easily  be  reduced  to  chromic  oxide  (see 
page  38)  by  boiling  its  solution  with  chlorhydric  acid  and 
alcohol,  or  by  adding  to  the  solution  sulphydric  acid  or  am- 
monic  sulphide.  Therefore,  in  all  solutions  to  which  these 
latter  reagents  have  been  added,  chromic  acid  is  changed  to 
chromic  oxide,  and  must  be  looked  for  among  the  metals. 

JBaric  Chloride  precipitates  chromic  acid  in  neutral 
solutions,  or  in  solutions  of  which  the  only  free  acid  is  acetic 
acid  (/.  e.,  solutions  which  have  been  made  alkaline  by  sodic 
hydrate,  and  acid  by  acetic  acid),  as  BARIC  CHROMATE, 
BaCrO4,  light  yellow.  Chromium  may  be  discovered  in  the  pre- 
cipitate of  baric  chromate  thus  obtained,  by  dissolving  some 
of  it  in  the  borax  bead.  The  green  color  of  chromic  oxide 
appears,  even  though  very  little  chromate  of  barium  was  con- 
tained in  the  precipitate. 

No  other  acid  which  precipitates  baric  chloride  gives  the  same 
color  in  the  borax  bead. 


SECTION  II. 

SULPHURIC  AND  SULPHUROUS  ACIDS. 

Acids  which  are  precipitated  by  baric  chloride  in  acid  solutions, 
either  immediately  or  after  oxidation. 

^Blowpipe  Reactions.  All  compounds  containing 
sulphur,  when  they  are  pulverized,  moistened,  and  mixed  with 
sodic  carbonate,  and  when  the  mixture  is  heated  on  charcoal, 
form  sodic  sulphide.  The  sodic  carbonate  must  be  heated 
until  it  soaks  into  the  charcoal,  then  if  the  portion  of  the  char- 
coal which  has  absorbed  it  is  dug  out  with  a  knife,  moistened 
with  water,  and  laid  on  a  piece  of  bright  silver,  the  presence 
of  sulphur  may  be  detected  by  the  appearance'of  a  black  stain 


SULPHURIC  AND  SULPHURO  US  A  CIDS.  65 

of  sulphide  of  silver.  If  no  silver  is  at  hand,  the  sodic  sul- 
phide can  be  extracted  by  soaking  with  water  the  portion  of 
the  charcoal  which  has  been  heated.  The  solution,  after  it  has 
been  filtered,  and  acidified  with  acetic  acid,  gives  a  precipitate 
of  SULPHIDE  OF  LEAD,  PbS,  black,  with  a  solution  of  ACETATE 

OF  LEAD. 

For  the  separation  of  sulphuric  and  sulphurous  acids,  see  Part 
III.  (123)  and  (124). 

SULPHURIC  ACID. 

H2SO4;  Na2SO4. 

Baric  Chloride  precipitates  compounds  of  sulphuric 
acid  in  neutral  or  acid  solutions  as  BARIC  SULPHATE,  BaSO4, 
white  powder.  Baric  sulphate  is  insoluble  in  chlorhydric,  nitric, 
sulphuric,  and  the  weaker  acids.  No  other  acid  gives  a  like 
precipitate  with  BaCl2,  in  acid  solutions.  COMPOUNDS  OF  LEAD 
AND  STRONTIUM  and  STANNIC  OXIDE  are  the  only  others  precipi- 
tated by  sulphuric  acid. 

Calcic  Chloride  in  concentrated  solutions  gives  a  pre- 
cipitate of  CALCIC  SULPHATE,  CaSO4,  but  calcic  sulphate  is 
soluble  in  380  parts  of  water. 

SULPHUROUS  ACID. 

H2SO8;  Na2SO8. 

Sulphurous  acid,  in  acid  solutions,  can  be  recognized  by  the 
smell  of  sulphurous  oxide  which  is  given  off.  Sulphurous 
oxide  is  evolved  from  acid  solutions  of  sulphites  when  they 
are  heated. 

JZaric  Chloride  precipitates  compounds  of  sulphurous 
acid  in  neutral  solutions  as  BARIC  SULPHITE,  BaSO3,  white 
powder.  Baric  sulphite  is  decomposed  and  dissolved  by  chlor- 
hydric and  nitric  acids.  POTASSIC  DICHROMATE  oxidizes  sul- 
phurous acid  in  a  solution,  made  acid  with  chlorhydric  acid,  to 
SULPHURIC  ACID,  2(H2CrO4)  +  3(H2SO8)  =  Cr2(SO*)8  +  5H2O. 
5 


66  PART  II. 

Sulphydric  Acid  decomposes  sulphurous  acid  in  acid 
solutions  with  precipitation  of  sulphur. 

2H2S  +  H2SO8  =  S3  +  3H2O. 

The  Iron  Reduction  Test.  POTASSIC  FERRICYAN- 
IDE  AND  FERRIC  CHLORIDE,  in  consequence  of  the  reduction 
of  the  ferric  chloride,  give  a  blue  precipitate  when  these  so- 
lutions are  mixed,  and  a  drop  of  the  mixture  is  held  on  the  end 
of  a  glass  rod  in  an  atmosphere  containing  sulphurous  acid. 
Sulphydric  acid  gives  the  same  reaction,  and  when  sulphydric 
acid  is  present  it  is  necessary  to  add  to  the  solution  sufficient 
plumbic  acetate  to  precipitate  it,  before  applying  this  test  for 
sulphurous  acid.  This  reaction  is  produced  equally  well  in  a 
solution  ;  but  there  are  many  other  reducing  agents  which  act 
in  the  same  manner,  but  which  do  not  take  the  gaseous  form, 
and  for  this  reason  it  is  best  to  add  chlorhydric  acid  to  the 
solution,  to  warm  it,  and  to  test  the  gas  evolved  for  sulphurous 
acid  with  potassic  ferricyanide  and  ferric  chloride. 

Sulphurous  acid  is  usually  recognized  by  its  smell,  or  ty  the 
above  test. 

SECTION  III. 
PHOSPHORIC,  BORACIC,  OXALIC,  AND  FLUORHYDRIC  ACIDS. 

Acids  which  are  precipitated  by  baric  chloride  in  neutral  solutions , 
but  which  are  not  precipitated  in  acid  solutions. 

For  the  tests  by  which  these  acids  can  be  most  easily  detected,  see 
Part  III.  (125-128). 

PHOSPHORIC    ACID. 

H3PO4;  Na2HPO4. 

All  the  Metals  of  Groups  II.,  III.,  IV.,  V.,  pre- 
cipitate compounds  of  phosphoric  acid  in  neutral  or  in  slightly 
alkaline  solutions  as  PHOSPHATES.  The  precipitate  is  soluble 


BORACIC  ACID.  67 

in  acids.  The  phosphoric  acid  cannot  usually  be  separated 
from  the  metal  by  treating  the  precipitate  with  ammonic 
hydrate,  and  from  most  metals  it  can  only  be  partially  separ- 
ated by  a  treatment  with  sodic  hydrate  or  carbonate.  MAG- 
NESIC  SULPHATE  solution,  to  which  a  solution  of  ammonic 
chloride  and  of  ammonic  hydrate  is  added,  precipitates  com- 
pounds of  phosphoric  acid  as  MAGNESIC  AMMONIC  PHOSPHATE, 
MgNH4PO4,  white  crystalline  powder,  soluble  in  acids  but  inso- 
luble in  ammonic  hydrate.  When  the  quantity  of  phosphoric 
acid  is  very  small,  the  precipitate  only  forms  after  some  time. 
This  is  the  best  reagent  for  phosphoric  acid  in  neutral  or  slightly 
alkaline  solutions. 

Molybdate  of  Ammonium  *  precipitates  compounds 
of  phosphoric  acid  in  an  acid  solution  as  a  yellow  crystalline 
phospho-molybdate  of  ammonium.  No  other  acid  gives  a  like 
precipitate.  The  best  reagent  for  phosphoric  acid  in  acid  solu- 
tions is  molybdate  of  ammonium,  when  there  are  bases  present 
which  are  precipitated  by  ammonic  hydrate. 

Sulphydric  acid  and  ferrocyanhydric  acid  also  precipitate 
ammonic  molybdate  solution.  (See  Part  III.,  127.) 

BORACIC  ACID. 

H8BO3 ;  Na2B4O7,  Borax. 

All  the  Metals  of  Groups  II.,  III.,  IV.,  V.  preci- 
pitate compounds  of  boracic  acid  in  neutral  or  slightly  alkaline 
solutions  as  BORATES.  The  precipitate  is  soluble  in  acids. 

The  precipitation  of  boracic  compounds  by  metals  of  these 
groups  is  mostly  incomplete,  and  the  acid  can  be  separated 
from  the  metal  in  almost  all  cases  by  treating  the  precipitate 
with  ammonia  or  with  sodic  hydrate  or  carbonate. 

Flame  Test.     Boracic  acid  in  dry  compounds  or  in  con- 

*  It  is  best  to  pour  a  few  drops  of  the  solution  to  be  tested  into  the  am- 
monic molybdate  solution,  instead  of  adding  the  ammonic  molybdate  to  the 
solution  to  be  tested. 


68  PART  II. 

centrated  solutions  imparts  a  beautiful  green  color  to  the  flame 
of  burning  alcohol. 

In  order  to  perform  the  test,  mix  the  dry  boracic  acid  com- 
pound with  a  few  drops  of  strong  sulphuric  acid  in  a  small 
evaporating  dish,  add  alcohol,  and  set  the  alcohol  on  fire.  Stir 
the  contents  of  the  dish  constantly  during  the  combustion  of 
the  alcohol,  and  observe  the  green  color,  when  the  alcohol  has 
mostly  burned  away. 

Turmeric  Paper  Test.*  Boracic  acid  compounds, 
when  their  solution  has  been  rendered  acid  by  chlorhydric 
acid,  turn  a  piece  of  turmeric  paper,  which  has  been  dipped  in 
the  solution  and  completely  dried,  brownish  red.  This  is  the 
usual  test  for  boracic  acid. 

OXALIC  ACID. 
H2C2O4;  KHC2O4- 

All  the  Metals  of  Groups  II.,  III.9  IV.,  F.  pre- 
cipitate salts  of  oxalic  acid  in  neutral  solutions  as  OXALATES, 
except  chromic  oxide  compounds.  The  precipitate  is  soluble 
in  acids,  and  in  the  case  of  many  metals  the  oxalic  acid  is 
removed  from  the  precipitate  by  ammonic  hydrate  and  by  sodic 
hydrate  and  carbonate. 

Sulphate  of  Calcium  precipitates  oxalic  acid  and  its 
salts  in  a  solution  to  which  sufficient  ammonic  hydrate  to 
render  the  solution  strongly  alkaline,  and  then  sufficient  acetic 
acid  to  render  the  reaction  acid,  have  been  added,  as  CALCIC 
OXALATE,  CaC2O4,  white  powder.  No  other  acid,  except  fluor- 
hydric  acid,  produces  a  precipitate  under  these  circumstances,  and 
calcic  fluoride  cannot  easily  be  mistaken  for  calcic  oxalate.  (See 
fluorhydric  acid,  below.) 

Sulphuric  Acid  (concentrated)  decomposes  dry  com- 
pounds, or  highly  concentrated  solutions  of  compounds  of 

*  If  iron  is  present  it  is  necessary  to  boil  the  solution  with  sodic  carbon- 
ate in  excess,  to  filter,  and  to  use  the  filtrate  for  the  turmeric  paper  test. 


FLUORHYDRIC  ACID. 


69 


oxalic  acid,  with  evolution  of  CARBONOUS  and  CARBONIC  OX- 
IDES.    Effervescence  takes  place.     H2C2O4  =  CO  +CO2  +  H2O. 

FLUORHYDRIC    ACID. 

HF  ;  CaF2 ;  NH4F. 

Fluorhydric  acid  cannot  be  present  in  acid  solutions  in  glass 
vessels. 

BARIUM,  STRONTIUM,  and  CALCIUM  salts  in  neutral  solu- 
tions precipitate  fluorhydric  acid  as  BARIUM,  STRONTIUM,  or 
CALCIUM  Fluoride,  BaF2,  SrF2,  CaF2.  The  two  former  are 
voluminous  white  precipitates.  Calcium  fluoride  is  a  gelatin- 
ous transparent  precipitate,  whose  formation  it  is  very  difficult 
to  observe. 

//  is  usually  unnecessary  to  test  for  fluorhydric  acid  except  in 
solid  substances. 

Sulphuric  Acid  (concentrated)  sets  fluorhydric  acid 
free  from  its  solid  compounds,  and  the  acid  may  be  recog- 
nized by  its  property  of  etching  glass.  To  perform  the  test, 
mix  the  pulverized  substance,  containing  fluorine,  with  strong 
sulphuric  acid  in  a  lead  cup,*  or  in  a  platinum  crucible.  Pre- 
pare a  piece  of  glass  by  melting  wax  on  it  and  pouring  off  all 
that  does  not  adhere,  leaving  a  thin  coating  of  wax  on  the  sur- 
face ;  scratch  lines  in  the  wax,  laying  the  surface  of  the  glass 
bare  ;  cover  the  vessel  containing  the  fluorine  compound  and 
sulphuric  acid  with  the  glass,  and  warm  gently.  After  fifteen 
minutes  warm  the  glass,  and  rub  off  the  wax  ;  the  surface  ex- 
posed by  the  scratches  will  be  etched  by  the  fluorhydric  acid. 

Fluorhydric  acid  must  always  be  removed,  if  present,  by  heat- 
ing with  strong  sulphuric  acid,  before  the  other  acids  of  Group  I., 
except  sulphuric  acid  and  carbonic  acid,  are  tested  for  in  a  sub- 
stance. 

*  A  lead  cup  may  be  made  by  hammering  up  the  end  of  a  piece  of  one- 
inch  lead  pipe  until  it  is  entirely  closed,  and  by  sawing  off  the  pipe  i$ 
inches  from  the  end.  Such  a  cup  may  be  warmed  on  a  sand-bath  suffi- 
ciently for  making  the  test  without  danger  of  melting  it. 


7o  PART  II. 

SECTION  IV. 
CARBONIC  ACID  AND  SILICIC  ACID. 

Acids  which  are  precipitated  by  baric  chloride  in  neutral  selu- 
tion,  but  which  are  set  free  by  acids,  and  which  cannot  be  present 
in  a' solution  that  has  been  evaporated  with  an  excess  of  an  acid. 

CARBONIC    ACID. 

Na2C03 ;  K2CO3. 

Carbonic  acid  can  only  be  present  in  considerable  quantity 
in  a  solution  which  has  an  alkaline  reaction. 

All  the  Metals  of  Groups  II.,  III.,  IV.,  V.  pre- 
cipitate alkaline  carbonates. 

Every  Acid  sets  free  CARBONIC  DIOXIDE,  CO2,  from  a 
solution  of  a  carbonate.  An  effervescence  or  the  formation  of 
bubbles  can  be  observed  when  an  acid  is  added  to  a  solution  of 
a  salt  of  carbonic  acid.  Carbonates  insoluble  in  water  are  also 
decomposed  by  a  free  acid,  carbonic  dioxide  being  evolved. 

Carbonic  dioxide  can  easily  be  recognized  by  the  white 
precipitate  of  CARBONATE  OF  CALCIUM,  CaCO3,  which  the  gas 
produces  in  a  drop  of  lime-water,*  held  in  it  on  the  end  of  a 
glass  rod. 

SILICIC    ACID. 

K2Si03. 

Silicic  acid  combined  with  bases  in  a  solution  is  set  free  by 
all  acids,  and  being  decomposed  into  water  and  silicic  oxide 
(H2SiO3  =  H2O  +  SiO2),  the  latter  often  separates  as  a  pre- 
cipitate ;  frequently,  however,  it  remains  for  months  in  a  solu- 
tion after  an  acid  has  been  added. 

*  The  lime-water  is  soon  destroyed  by  the  absorption  of  carbonic  di- 
oxide from  the  air,  and  before  using  it  to  test  for  carbonic  acid,  the  stu- 
dent should  assure  himself  that  a  drop  of  it  gives  a  precipitate  with  the 
gas  evolved  from  sodic  carbonate,  to  which  chlorhydric  has  been  added. 


SILICIC  ACID.  7I 

When  silicic  oxide,  or  silica,  has  been  separated  from  its  com- 
bination with  a  base  by  the  addition  of  an  acid,  and  the  solution 
has  been  evaporated  completely  to  dryness,  the  silica  remains 
perfectly  insoluble  when  the  dry  mass  is  treated  with  water  or 
acids  to  dissolve  the  bases.  This  is  the  characteristic  test  for 
silicic  acid. 

When  silicic  acid  is  present  in  a  solution  it  is  recognized 
and  separated  by  the  above  process  before  any  other  tests  are 
performed. 

Many  solid  compounds  of  silicic  acid  are  not  acted  upon  by 
acids,  and  can  only  be  brought  into  solution  by  the  process 
described,  Part  III.,  XXIII. 

The  following  properties  of  silicic  acid  must  be  considered, 
in  order  to  determine  whether  a  body,  which  has  been  left  in- 
soluble after  a  treatment  with  an  acid,  is  silica,  or  some  other 
compound,  which  is  likewise  insoluble  : 

Silicic  acid  is  precipitated  from  its  solutions  by  acids  in  the 
gelatinous  form,  or  in  the  form  of  amorphous  white  flakes. 
Silica  which  has  been  dried  always  has  the  latter  form. 

Silica,  after  it  has  been  fused  with  four  parts  of  a  mixture 
of  sodic  carbonate  and  potassic  carbonate,  forms  a  glass,  which 
is  entirely  soluble  in  water.  When  the  solution  thus  obtained 
is  evaporated  with  an  excess  of  nitric  acid,  and  the  dry  mass 
is  treated  with  water,  no  metal  should  go  into  solution  which 
gives  a  precipitate  with  sodic  carbonate. 


PART  77. 


GROUP  II. 

CHLORHYDRIC,  BROMHYDRIC,  IODHYDRIC,  CY- 
ANHYDRIC,  FERROCYANHYDRIC,  FERRICYAN- 
HYDRIC,  AND  SULPHYDRIC  ACIDS. 

Acids  which  are  not  precipitated  by  baric  chloride,  but  which 
are  precipitated  by  argentic  nitrate  in  nitric  acid  solution. 

SECTION  I. 

CHLORHYDRIC,  BROMHYDRIC,  IODHYDRIC,  AND  CYANHYDRIC 

ACIDS. 

Acids  which  give  with  argentic  nitrate  a  white  or  light  yellow 
flocculent  precipitate,  insoluble  in  dilute  nitric  acid,  and  are  not 
precipitated  by  salts  of  iron  in  acid  solution. 

CHLORHYDRIC    ACID. 

HC1;  NaCl. 

Plumbic,  Mercurous,  and  Argentic  Salts  are 

the  only  compounds  which  give,  with  chlorhydric  acid,  pre- 
cipitates insoluble  in  nitric  acid. 

Sulphuric  Acid  (concentrated)  sets  free  chlorhydric 
acid  from  its  compounds.  Chlorhydric  acid  gas  precipitates  a 
drop  of  argentic  nitrate,  held  on  the  end  of  a  glass  rod  in  an 
atmosphere  containing  it,  as  ARGENTIC  CHLORIDE,  AgCl,  white 
flakes.  Compounds  containing  cyanhydric,  chloric,  and  hypo- 
chlorous  acids,  produce  the  same  reaction.  Chlorhydric  acid 
gas  does  not  bleach  indigo  solution.  See,  however,  Part  III. 
(48). 


BROMHYDRIC  AND  IODHYDRIC  ACIDS. 


73 


Argentic  Nitrate  precipitates  solutions  containing 
chlorhydric  acid  as  ARGENTIC  CHLORIDE,  AgCl,  white  flakes, 
turning  purple  on  exposure  to  light,  settling  quickly  after  they 
have  been  shaken.  Argentic  chloride  is  soluble  in  ammonic 
hydrate,  and  completely  insoluble  in  boiling  nitric  acid  (con- 
centrated). No  other  compound  of  silver  except  the  ferro-  and 
ferricyanide  remains  undissolved  after  this  treatment.  When 
ferro-  or  ferricyanhydric  acid  is  present  in  a  solution  containing 
chlorhydric  acid,  follow  the  directions  given,  Part  III.  (130) 
(a),  before  applying  the  argentic  nitrate  test. 

This  is  the  usual  test  for  chlorhydric  acid. 

BROMHYDRIC    ACID. 

KBr. 

Nearly  all  bromides  are  soluble  in  water  ;  bromide  of  lead, 
however,  dissolves  very  sparingly,  and  the  mercurous  and 
argentic  salts  are  quite  insoluble. 

Argentic  Nitrate  precipitates  from  solutions  of  bro- 
mides the  ARGENTIC  BROMIDE,  AgBr,  white,  closely  resembling 
the  chloride,  but  more  difficultly  soluble  in  ammonia  ;  insolu- 
ble in  hot  nitric  acid. 

Bromides  are  decomposed  by  chlorine,  hypochlorites,  strong 
sulphuric  and  nitric  acids,  bromine  being  set  free,  which  im- 
parts a  yellow  or  yellowish-red  color  to  the  liquid.  On  shak- 
ing the  tube  with  a  little  ether  or  carbonic  disulphide  the 
bromine  will  be  dissolved  and  impart  to  the  solvent  a  yellow 
or  reddish-brown  color. 

IODHYDRIC    ACID. 

KI. 

The  iodides  resemble  very  much  the  corresponding  chlorides 
and  bromides. 
Argentic  Nitrate  produces  in  solutions  of  iodides  a 


74  PART  II. 

yellowish-white  precipitate  of  ARGENTIC  IODIDE,  Agl,  very 
slightly  soluble  in  ammonia,  but  soluble  in  boiling  concen- 
trated nitric  acid. 

JVLeTCUTOUS  NitTate  precipitates  yellowish-green  MER- 

CUROUS  IODIDE,  Hg2I2. 

jMLevcwT-ic  Chloride  precipitates  brilliant  scarlet  MER- 
CURIC IODIDE,  HgI2. 

Chlorine  or  bromine  water  liberates  iodine  from  iodides  ;  on 
shaking  with  carbonic  disulphide  the  iodine  is  concentrated  in 
this  liquid,  forming  a  violet-colored  solution. 

Free  iodine  imparts  a  blue  color  to  starch  paste.  For  this 
purpose  dilute  starch  paste  is  added  to  a  solution  of  an  iodide, 
and  then  chlorine  water,  nitric  acid,  or  sulphuric  acid  carefully 
added  to  liberate  the  iodine.  If  too  much  chlorine  is  added 
chloride  of  iodine  is  formed,  which  prevents  the  formation  of 
the  blue  color.  The  chloride  of  iodine  is  also  produced  by 
nitric  acid  in  presence  of  considerable  amount  of  chlorides. 


CYANHYDRIC  ACID. 

HCy;  KCy. 

The  reactions  of  different  classes  of  cyanhydric  acid  com- 
pounds must  be  considered  separately. 

Soluble  Simple  Cyanides.  Cyanides  of  metals  of 
Groups  I.,  II.,  and  III.  are  soluble  in  water.  Cyanhydric  acid 
is  set  free  from  their  solutions  by  even  the  feeblest  acids  (acetic 
and  carbonic).  (MERCURIC  CYANIDE  is  soluble  in  water,  but 
is  not  decomposed  by  alkalies  nor  by  acids,  except  by  sul- 
phydric  acid  ;  and  the  tests  described  for  cyanhydric  acid 
cannot  be  applied  to  it.  Sulphydric  acid  precipitates  mercu- 
ric sulphide,  and  sets  cyanhydric  acid  free. ) 

Insoluble  Simple  Cyanides.  Cyanides  of  metals 
of  Groups  IV.  and  V.,  except  mercuric  cyanide,  are  insoluble 
in  water,  and  the  cyanides  of  metals  of  Group  V.  are  not  de- 


CYANHYDRIC  ACID. 


75 


composed,  or  are  decomposed  with  great  difficulty  by  acids. 
The  insoluble  cyanides  dissolve  readily  in  potassic  cyanide,  and 
the  ordinary  tests  for  metals  cannot  be  used  with  such  solutions. 
The  cyanides  can  be  precipitated  from  these  solutions  by  the 
addition  of  an  acid,  with  some  exceptions,  the  two  most  re- 
markable of  which  are  described  separately.  (See  Ferro-  and 
Ferricyanhydric  Acids.) 

Free  Cyanhydric  Acid  can  be  recognized  by  its 
smell,  which  is  like  that  of  bitter  almonds.  (The  acid  is  very 
poisonous,  and  the  fumes  arising  from  a  solution  containing  a 
considerable  quantity  of  it  should  be  inhaled  with  caution.) 

Argentic  Nitrate  precipitates  soluble  compounds  of 
cyanhydric  acid  in  an  acid  solution  as  ARGENTIC  CYANIDE, 
AgCy,  white  flakes,  which  do  not  settle  so  readily,  when  shaken, 
as  argentic  chloride,  and  which  do  not  turn  purple  quickly  in  the 
light.  Argentic  cyanide  is  wholly  decomposed  and  dissolved 
by  boiling  a  few  minutes  with  strong  nitric  acid.  It  is  also 
decomposed,  and  cyanhydric  acid  goes  into  solution,  when  it  is 
digested  with  dilute  chlorhydric  acid  in  contact  with  metallic 
zinc. 

Prussian  Blue  Test.  When  a  feebly  acid  solution 
containing  cyanhydric  acid  is  mixed  with  several  drops  of 
ferrous  sulphate  solution,  and  with  a  drop  of  ferric  chloride 
solution,  and  sodic  hydrate  is  added  until  a  precipitate  forms, 
and  the  mixture  is  warmed  for  a  minute,  and  then  acidified 
with  dilute  chlorhydric  acid,  a  blue  precipitate,  or  more  frequently 
a  blue  coloration,  appears,  either  immediately  or  after  the  addi- 
tion of  a  drop  of  ferric  chloride. 

When  ferro-  or  ferricyanhydric  acid  or  both  acids  are  present 
(see  the  following  section),  they  must  be  removed  from  the 
solution  before  the  Prussian-blue  test  can  be  applied.  To  this 
end  add  to  a  small  quantity  of  the  solution  an  equal  bulk  of 
dilute  sulphuric  acid,  and  dilute  with  a  considerable  quantity 
of  water  ;  add  ferric  chloride  or  ferrous  sulphate,  or  both  to- 
gether, according  as  ferro-  or  ferricyanhydric  acid  or  both 


76  PART  II. 

acids  are  present,  and  then  add  baric  chloride  until  the  blue 
precipitate  appears  of  a  much  lighter  shade  ;  shake  thoroughly, 
and  allow  the  precipitate  to  settle  for  a  few  minutes,  and  filter. 
If  only  the  first  few  drops  run  through  the  filter  blue,  they 
should  be  thrown  away  and  the  remainder  of  the  filtrate  taken. 
If  no  clear  filtrate  can  be  obtained,  add  to  the  filtrate  a  little 
baric  chloride  and  filter  again.  The  filtrate  is  to  be  tested  as 
above  by  the  addition  of  sodic  hydrate,  and  afterwards  of  an 
acid  for  cyanhydric  acid.  A  sufficiently  capacious  flask  must 
be  chosen  for  the  operation.  The  only  object  in  adding  baric 
chloride  is  to  facilitate  the  filtration  from  the  blue  precipitate. 
Cyanhydric  acid  is  the  only  acid  which  gives  this  reaction  under 
these  circumstances. 

SECTION  II.   . 
FERROCYANHYDRIC  AND  FERRICYANHYDRIC  ACIDS. 

Acids  which  give  with  argentic  nitrate  colored  precipitates, 
which  are  not  wholly  destroyed  on  boiling  with  strong  nitric  acid, 
and  which  are  precipitated  by  ferrous  or  ferric  salts,  and  by  cu- 
pric  salts  in  dilute  acid  solutions. 

FERROCYANHYDRIC    ACID. 

H4(FeCy6)  ;  K4(FeCy6). 

Ferrous  Sulphate  precipitates  ferrocyanhydric  acid 
compounds  in  acid  solutions  as  POTASSIC  FERROUS  FERROCY- 
ANIDE,  K2Fe(FeCy6),  bluish  white  precipitate,  which  quickly 
turns  dark  blue  through  oxidation  by  the  air. 

Ferric  Chloride  precipitates  ferrocyanhydric  acid  com- 
pounds in  acid  solution  as  PRUSSIAN  BLUE,  Fe4(FeCy6)3,  deep 
blue. 

Ferrocyanhydric  acid  is  the  only  acid  which  gives  this  reaction. 

Cupric  Sulphate  precipitates  ferrocyanhydric  acid  com- 
pounds in  acid  solution  as  CUPRIC  FERROCYANIDE  (Cu2FeCy6), 
brownish-red  powder. 


SULPH  YDRIC  A  CID.  j  j 

The  metals  are  left  as  oxides,  and  the  ferrocyanogen  is  dis- 
solved as  sodic  ferrocyanide,  when  these  precipitates  are  di- 
gested with  sodic  hydrate. 

FERRICYANHYDRIC    ACID. 

H6(Fe2Cy12)  ;  K6(Fe2Cy12). 

Ferrous  Sulphate  precipitates  compounds  of  ferricy- 
anhydric  acid  in  acid  solution  as  TURNBULL'S  BLUE,  Fe3(Fe2 
CyM),  deep  blue. 

Ferricyanhydric  acid  is  the  only  acid  which  gives  this  reaction. 

Ferric  Chloride  does  not  precipitate  compounds  of  fer- 
ricyanhydric  acid  in  acid  solution.  The  color  of  the  solution 
is  deepened. 

Cupric  Sulphate  precipitates  compounds  of  ferricyan- 
hydric  acid  in  acid  solution  as  CUPRIC  FERRICYANIDE,  yellow- 
ish-green powder. 

The  metals  are  left  as  oxides,  and  the  cyanogen  is  dissolved 
as  sodic  ferricyanide  when  these  precipitates  are  treated  with 
sodic  hydrate. 

SECTION  III. 

SULPHYDRIC   ACID. 

An  acid  which  gives  a  black  precipitate  with  salts  of  lead,  silver, 
copper,  and  many  others  in  an  acid  solution. 

No  other  acid  gives  a  precipitate  of  the  same  color  with  these 
metals. 

SULPHYDRIC    ACID. 

H2S  ;  (NH4)2S. 

Sulphydric  Acid  is  set  free  from  its  solutions  by  all 
other  acids  except  carbonic  and  cyanhydric  acid,  and  it  can 
be  recognized  by  its  smell.  The  sulphydric  acid  gas  is  given 
off  with  effervescence  when  the  solution  is  concentrated. 


7 8  PART  II. 

The  Metals  of  Groups  I.  and  II.  form  with 
sulphydric  acid  soluble  sulphides,  which  have  an  alkaline  re- 
action. 

The  Metals  of  Group  IV. 9  when  the  acid  with  which 
they  are  combined  is  neutralized,  form  with  sulphydric  acid 
INSOLUBLE  SULPHIDES,  which,  with  the  exception  of  the  sul- 
phides of  cobalt  and  nickel,  are  dissolved  by  cold  dilute  chlor- 
hydric  acid,  with  evolution  of  sulphydric  acid. 

Metals  of  Groups  V.  and  VI.  form  with  sulphydric 
acid  insoluble  sulphides,  which  are  not  decomposed  by  dilute 
acids.  (See  also  Mercury,  page  48  and  page  50.) 

Lead-Paper  Test.  A  piece  of  paper  moistened  with 
plumbic  acetate,  and  held  over  a  solution  from  which  sulphy- 
dric acid  is  set  free  by  the  addition  of  a  stronger  acid,  is  black- 
ened. No  other  acid  gives  this  reaction. 


NITRIC  AND  CHLORIC  ACIDS. 


79 


GROUP  III. 

NITRIC,  CHLORIC,  AND  ACETIC  ACIDS. 

Acids  which  are  not  precipitated  by  any  metal. 

SECTION  I. 

NITRIC  AND  CHLORIC  ACIDS. 
Acids  which  deflagrate  when  tested  with  the  blowpipe  on  charcoal. 

NITRIC    ACID. 

HNO8;  NaNOs. 

Nitric  Acid9  when  concentrated,  is  readily  decomposed 
when  heated  with  copper  turnings,  and  red  fumes  of  NITRIC 
PEROXIDE,  NO2,  are  given  off.  The  reaction  can  be  obtained 
with  a  moderately  dilute  solution  by  adding  to  it  concentrated 
sulphuric  acid.  No  reaction  is  obtained  with  very  dilute  so- 
lutions. 

J?errous  Sulphate  Test.  Add  a  few  drops  of  a  solu- 
tion containing  nitric  acid  to  concentrated  sulphuric  acid  in  a 
test-tube,  and  pour  upon  this  solution  a  layer  of  cold  ferrous 
sulphate  solution.  A  brown  or  red  color  appears  at  the  line  of 
separation  of  the  two  solutions,  arising  from  the  absorption  of 
nitrous  gases  by  the  ferrous  sulphate. 

This  is  the  characteristic  test  for  nitric  acid. 

CHLORIC    ACID. 
KC108. 

Sulphuric  Acid  (concentrated).    When  a  small  quan- 


80  PART  II. 

tity  of  a  solid  chlorate,  or  a  very  concentrated  solution  con- 
taining a  chloric  acid  compound,  is  added  to  strong  sulphuric 
acid,  and  heat  is  applied,  a  peculiar  yellow  gas  (oxides  of 
chlorine)  is  evolved,  which  has  a  characteristic  suffocating 
odor,  which  precipitates  ARGENTIC  CHLORIDE  in  a  drop  of  an 
argentic  nitrate  solution,  and  which  bleaches  a  drop  of  an  in- 
digo solution  when  these  reagents  are  held  on  the  end  of  a 
glass  rod  in  an  atmosphere  containing  the  gas. 

This  is  the  characteristic  test  for  chloric  acid. 

Hypochlorous  Acid,  gives  the  same  reactions  as  chloric 
acid,  but  that  acid  is  easily  set  free  and  evolved  from  its  solu- 
tion by  dilute  sulphuric  acid,  while  chloric  acid  is  not,  and 
moreover  it  is  usually  present  only  in  alkaline  solutions. 

Chlorhydric  Acid  in  the  presence  of  an  oxidizing 
agent  gives  a  similar  reaction,  but  the  yellow  gas  evolved 
(chlorine)  is  much  less  intense  in  color,  and  has  a  different 
odor.  It  is,  however,  very  difficult  to  distinguish  between  the 
reaction  given  by  chlorine  in  such  a  case  and  that  given  by 
chloric  acid  compounds. 

SECTION  II. 

ACETIC  ACID. 

An  acid  which  does  not  deflagrate  on  charcoal. 

ACETIC    ACID. 

HC2H802 ;  NaC2H802. 

The  Strong  Mineral  Acids  set  acetic  acid  free 
from  its  combinations. 

Acetic  acid  can  be  recognized  by  the  odor  of  vinegar  pecu- 
liar to  it. 

Sulphuric  Acid  Test.  When  an  equal  bulk  of  alco- 
hol is  added  to  strong  sulphuric  acid,  and  a  small  quantity  of 
a  solution  containing  a  compound  of  acetic  acid  is  added,  and 


ACETIC  ACID.  8 1 

the  mixture  is  heated,  a  characteristic  odor  of  acetic  ether  is 
given  off. 

In  case  gases  are  given  off,  which  make  it  difficult  to  recog- 
nize the  odor  of  acetic  ether,  it  is  advisable  to  provide  the  test- 
tube  in  which  the  reaction  is  performed  with  a  tube  for  distil- 
lation,* and  to  distil  a  small  quantity  of  the  alcohol  into  an- 
other test-tube,  to  mix  the  distillate  with  water,  to  neutralize  it 
with  sodic  carbonate,  and  to  warm  it ;  the  odor  of  acetic  ether 
can  then  be  recognized  in  the  liquid  which  was  distilled. 

This  is  the  characteristic  test  for  acetic  acid. 

Argentic  and  Mercurous  Nitrates  precipitate 
concentrated  neutral  solutions  of  acetic  acid  compounds  as 

ARGENTIC     AND    MERCUROUS    ACETATES,    AgC2H3O2  and  Hg2 

(C2H3O2)2,  white  crystalline  scales.  The  precipitates  are  soluble 
in  dilute  nitric  acid,  and  also  in  a  large  quantity  of  water. 

*  Bend  a  3-16  inch  tube,  of  about  one  foot  in  length,  at  an  angle  of 
about  80°,  so  that  one  arm  shall  only  be  i^  inches  long.  Fit  a  cork  to 
the  test-tube,  and  insert  the  bent  tube  in  a  hole  bored  through  the  cork 
with  a  round  file, 

6 


PART   III. 


PRELIMINARY  TESTS  WITH  NON-METALLIC 
SOLIDS. 

EXAMINATION  IN  A  CLOSED  TUBE. 

USE  a  piece  of  hard  glass  tubing  three-eighths  of  an  inch 
in  diameter,  closed  at  one  end  (see  page  26)  for  this  examina- 
tion. Introduce  the  substance,  pulverized  or  in  small  pieces, 
into  the  tube,  wipe  the  inside  of  the  tube  if  necessary  with  a 
bit  of  rolled  filter-paper,  and  heat  the  substance,  gently  at  first, 
but  eventually  to  the  highest  temperature  attainable  with  the 
flame  of  a  Bunsen's  lamp  or  with  the  blowpipe  flame.  Ob- 
serve carefully  the  changes  which  occur. 
No  Change.  The  substance  contains  no  organic  matter, 

(1)  no  readily  fusible  body,  no  readily  volatile  body,  and 
no  water. 

Pass  to  the  Examination  on  Charcoal  (page  85). 
Wat er*     Substances  containing  water  (usually  water  of  crys- 

(2)  tallization)  deposit  a  film  of  moisture  in  the  upper 
part  of  the  tube  when  they  are  heated.     If  the  water 
colors  turmeric  paper  brown,  AMMONIA  is  present. 

Organic  Matter*     Substances  containing  organic  matter 

(3)  blacken  and  give  off  gases  when  they  are  heated. 

Should  the  substance  contain  organic  matter  it  must 
be  burnt,  until  the  organic  matter  is  completely  de- 
stroyed,* by  heating  with  the  lamp  or  blowpipe,  on 

*  In  some  special  cases,  as  in*  examinations  for  mercury  and  arsenic, 
other  processes  of  analysis  must  be  employed,  for  which  larger  works  must 
be  consulted. 

82 


EXAMINA  TION  OF  SOLIDS.  83 

platinum  foil  or  on  a  bit  of  porcelain,  or  in  a  porcelain 
dish  or  crucible,  before  further  analysis,  commencing 
with  the  examination  on  charcoal  (page  85),  is  pro- 
ceeded with. 

A  GAS  IS  GIVEN  OFF. 

May  be  recognized  by  its  property  of  rekindling 
(4)     a  glimmering  match  held  in  the  tube. 

PEROXIDES,  NITRATES,  and  CHLORATES  evolve  oxy- 
gen. 

Nitrates  and  chlorates  also  deflagrate  on  charcoal. 
See  (14). 

Sulphurous  Oxide,  SO2,  can  be  recognized  by  its  smell. 
(3)         Some    SULPHATES    of    higher    metals,    and    many 
SULPHITES,   evolve   sulphurous  oxide   when  they  are 
heated.* 

SulpTiydric  Acid,  H2S,  can  be  recognized  by  its  smell 
(0)     and  by  its  property  of  blackening  lead  paper.    See 
(45). 

Some  alkaline  SULPHIDES,  containing  water,  evolve 
sulphydric  acid  when  they  are  heated. 
Carbonic  Dioxide,  CO2,  can  be  recognized  by  its  prop- 

(7)  erty  of  extinguishing  a  lighted  or  glimmering  match 
held  in  the  tube.     See  also  (41). 

Some  CARBONATES  lose  carbonic  dioxide  when  they 
are  heated. 
Hyponitric  Oxide,  NO2,  appears  as  red  fumes. 

(8)  NITRATES  of  the  higher  metals  evolve  hyponitric 
oxide  when  they  are  heated. 

Ammonia,  NH3,  can  be  recognized  by  its  smell,  and  by  its 

*  Many  sulphides  of  higher  metals  give  off  sulphur  in  the  form  of  sul- 
phurous oxide  when  they  are  roasted  with .  access  of  air.  The  sulphides, 
finely  pulverized,  may  be  heated  red-hot  in  a  tube,  open  at  both  ends,  and 
held  in  an  inclined  position  to  favor  the  draught,  and  the  sulphurous  oxide 
may  be  detected  by  its  smell  at  the  upper  end  of  the  tube. 


84  PART  III. 

(9)  property  of  turning  moist  turmeric  paper  brmvn.    Salts 
of  ammonia,  in   the  presence  of  alkalies,  and  some 
organic  substances,  evolve  ammonia  when   they   are 
heated. 

[Better  tests  for  these  bodies,  with  the  exception  of  oxygen, 
are  given  in  the  following  pages,  since  it  is  often  difficult  to 
observe  the  formation  of  a  gas  in  a  small  tube  ;  the  phenomena 
described  above  should,  however,  be  looked  for  when  sub- 
stances are  heated  in  a  closed  tube.] 

A  SUBLIMATE  FORMS. 

An  opinion  may  be  formed  of  the  volatility  of  the  sublimate, 
according  to  the  distance  from  the  heated  part  of  the  tube  at 
which  it  is  deposited. 
Sulphur   sublimes    easily  and   solidifies    in    reddish-brown 

(10)  drops,    which  become  yellow   or  yellmrish-brown  on 
cooling. 

Some  METALLIC    SULPHIDES  give  off  a  portion  of 
their  sulphur  when  they  are  heated. 
Ammonic  Salts  form  white  sublimates.     Touch  the  sub- 

(11)  limate  with  a  drop  of  sodic  hydrate,  or  with  a  bit  of 
paper  moistened  with  sodic  hydrate,  and  if  the  smell 
of  ammonia  is  given  off  it  consists  of  an  ammonic  salt. 

Mercury.     Metallic  mercury  sublimes  as  &  grey  film,  which 

(12)  augments  to  form  globules  when  the  quantity  of  mer- 
cury is  large. 

MERCURIC  SULPHIDE,  HgS,  gives  a  black  sublimate, 
which  becomes  ra/when  it  is  rubbed. 

MERCUROUS    CHLORIDE,    Hg2Cl2,   and    MERCURIC 
CHLORIDE,  HgCl2,  give  a  white  sublimate,  which  turns 
black  when  it  is  moistened  with  ammonic  sulphide  so- 
lution. 
Arsenic.     Metallic  arsenic  sublimes  and  deposits  itself  as  a 

(13)  brilliant  black  metallic  ring  in  the  tube. 


EXAMINA  TION  ON  CHARCOAL.  85 

ARSENIOUS  OXIDE,  As2O8,  forms  a  white  crystalline 
sublimate,  which  turns  yellow  when  it  is  moistened 
with  sulphydric  acid  solution. 

ARSENIOUS  SULPHIDE,  As2S3,  forms  a  sublimate, 
which  is  reddish  yellow  when  hot,  and  yelloiv  when 
cold.  It  is  somewhat  less  volatile  than  sulphur. 


RECAPITULATION  (10)  TO  (13)  SUBLIMATES. 

The  substance  is  heated  in  a  closed  tube. 

WHITE  SUBLIMATE — ammonic  salts  (11);  mercurous  chloride, 

Hg2Cl2,  and  mercuric  chloride,  HgCl2  (12)  ;  and  arsenious 

oxide,  As2O3  (13). 
YELLOW  SUBLIMATE — sulphur  (10)  ;  and  arsenious  sulphide, 

As2S3  (13). 

BROWN  SUBLIMATE  (while  hot) — sulphur  (10). 
REDDISH-YELLOW  SUBLIMATE  (while  hot) — arsenious  sulphide, 

As2S3  (13). 

GRAY  METALLIC  SUBLIMATE — mercury  (12). 
BLACK  SUBLIMATE — arsenic    (13)  ;  and  mercuric   sulphide, 

HgS,  red  when  rubbed  (12). 


EXAMINATION  ON  CHARCOAL. 

Hollow  out  a  small  cavity  in  a  piece  of  charcoal  (see  page 
25),  and  heat  a  portion  of  the  solid  substance  with  the  blow- 
pipe flame. 
titrates  and  Chlorates  enter  into  a  vivid  combus- 

tion,  called  deflagration,   when  they  are  heated  on 

charcoal. 


86  PART  III. 

Potassium  and  Sodium  Salts  melt,  and  some  of 
(15)    them  are  imbibed  by  the  pores  of  the  charcoal  when 

they  are  heated. 

Compounds  of  the  Metals  of  Groups  II.  and 
(16)  III.,  also  Zinc  Compounds  and  Silicic 
Oxide,  remain  as  a  white  infusible  mass  on  the  char- 
coal after  heating.  Frequently,  when  heat  is  first  ap- 
plied, they  melt  in  their  water  of  crystallization,  and 
afterwards  become  solid. 

ALUMINIC  OXIDE  becomes  blue,  and  ZINC  OXIDE  be- 
comes green  when  they  are  moistened  with  cobaltic 
nitrate  and  heated  in  the  oxidizing  flame. 
Salts  of  the  Metals  of  Groups  IV.  and  V.  leave  a 

(17)  dark-colored  residue  when  they  are  heated  on  char- 
coal.    The  oxides  of  these  metals  generally  assume  a 
darker  color  when  they  are  heated.    Exceptions  :  Zinc 
and  mercury. 

Salts  of  Ammonia  and  Mercury,  also  Com- 

(18)  pounds  of  Arsenic  and  Antimony,  which 
do  not  contain  another  metal,  volatilize   completely 
when  they  are  heated  on  charcoal. 

Gold  and  Silver  Compounds,  also  Oxides  of 

(19)  Lead  and  Bismuth,  give  bright  metallic  globules 
when  they  are  heated  on  charcoal. 


FUSION  WITH  SODIC  CARBONATE. 

(20)  When  metals  of  Groups  IV.,  V.,  and  VI.  appear  to 
be  present  (see  17),  mix  a  small  quantity  of  the  pul- 
verized substance  with  two  or  three  times  its  bulk  of 
sodic  carbonate  in  the  palm  of  the  hand,  moisten  with 
water,  and  form  the  mixture  by  working  it  with  a  knife- 
blade  into  a  ball  the  size  of  a  pea.  Place  the  ball  in 
a  cavity  scooped  out  of  a  piece  of  charcoal,  and  heat 


EX  AM  IN  A  TION  ON  CHARCOAL.  gy 

with  the  inner  blowpipe  flame  until  almost  all  of  the 
carbonate  of  soda  has  been  imbibed  by  the  charcoal. 
Many  metals  are  reduced  and  appear  as  metallic 
globules  in  the  cavity  of  the  charcoal,  and  those  which 
are  volatile  deposit  an  incrustation  of  their  oxides  on 
the  charcoal.  This  incrustation  is  to  be  looked  for  at 
a  greater  or  less  distance  from  the  cavity,  according 
to  the  volatility  of  the  metal,  and  always  in  the  direc- 
tion in  which  the  metallic  vapors  are  blown  by  the 
flame. 

The  physical  and  chemical  properties  of  the  glob- 
ules and  the  color  of  the  incrustations  afford  means  of 
recognizing   several   metals,   usually,    however,    only 
when  they  are  not  associated  with  others. 
Tron,  Cobalt,  Nickel,  and   Manganese    Com- 

(21)  pounds  give  neither  globule  nor  incrustation. 
Gold,  Silver,  and  Copper  Compounds  give  mallea- 

(22)  ble  globules,  which  can  be  distinguished  by  the  re- 
spective colors  of  the  metals.     They  give  no  incrusta- 
tion. 

Zinc  Compounds  give  no  globules,  but  a  white  incrusta- 

(23)  tion,  ZnO,  near  the  spot  heated.     The  incrustation  is 
yellow  while  hot.     It  is  not  volatile  in  the  oxidizing 
flame.     It  becomes  green  when  it  is  moistened  with 
nitrate  of  cobalt,  and  heated  in  the  oxidizing  flame. 

Tin  Compounds  give  very  ductile  white  globules. 

(24)  The  incrustation,  SnO2,  produced  by  tin  compounds 
is  dirty  yellow  when  hot,  and  lighter  when  cold.     It  is 
deposited  in  the  immediate  vicinity  of  the  cavity,  and 
it  is  very  difficult  to  distinguish  it  from  the  ash  of  the 
charcoal. 

Lead  Compounds  give  very  ductile  globules. 

(25)  The  incrustation,  PbO,  is  bright  yellow  when  hot, 
and  pale yellow  when  cold.     It  is  deposited  at  a  greater 
distance  than  SnO2  from  the  cavity.     When  the  blow- 


88  PART  III. 

pipe  flame  is  directed  upon  the  incrustation  of  PbO 
it  vanishes,  and  the  flame  is  colored  blue. 
Bismuth  Compounds  give  brittle  globules.     The  incrus- 

(26)  tation,  Bi2O3,  is   orange  yellow  when  hot,  and  bright 
yellow  when   cold.     It  vanishes  when   the   blowpipe 
flame  is  directed  upon  it,  but  it  does  not  impart  a  blue 
color  to  the  flame. 

Cadmium,  Compounds  give  no  globules,  but  a  yellow  to 

(27)  reddish-brown  incrustation,  very  different  in  color  from 
that  of  any  other  metal. 

Antimony  Compounds  give  brittle  globules,  but  metallic 

(28)  antimony  is  so  volatile  that  frequently  these  are  driven 
off  by  the  heat  required  for  their  reduction.     Some- 
times fumes,  arising  from*  the  vapor  of  antimony,  are 
visible.     The  incrustation,  Sb2O3,  is  white.     It  is  de- 
posited at  a  greater  distance  than  PbO  from  the  cavity, 
and  it  can  easily  be  driven  from  one  place  to  another 
on  the  charcoal  by  the  heat  of  the  blowpipe  flame. 

Arsenic  Compounds  give  no  globules,  but  a  character- 
(20)  istic  garlic  odor.     The  incrustation,  As2O3,  is  white, 

and  it  is  still  more  volatile  than  Sb2O3. 
When  the  compound  contains  several  metals  that  can  be  re- 
duced, they  alloy  with  each  other,  and  it  is  usually  impossible 
to  recognize  the  metals  in  the  presence  of  each  other  by  their 
physical  properties  ;  also,  the  incrustation  given  by  one  metal 
frequently  obscures  that  given  by  another. 

If  a  sufficient  quantity  of  the  metal  can  be  easily  reduced, 

it  is  always  advisable  to  treat  it  with  solvents  in  the  manner  to 

be  described  under  metals.     (See  page  98.) 

Sulphur.     The  following  modification  of  the  fusion  with 

(30)  carbonate  of  sodium  on  charcoal  is  a  valuable  test  to 

discover  sulphur  in  baric  sulphate  and  in  sulphides. 

If  a  piece    of  the  charcoal  which  has  imbibed  the 

soda  is  moistened  and  laid  on  a  silver  coin,  a  black 

stain  appears,  if  the  coin  is  washed  after  a  few  minutes, 


EXAMINA  TION  ON  CHARCOAL.  89 

in  case  sulphur  is  present.  The  charcoal  may  also  be 
pulverized  and  treated  with  water,  and  if  the  solution, 
after  being  filtered,  gives  a  black  precipitate  with 
plumbic  acetate  solution,  sulphur  is  present.  For  this 
test  the  flame  of  a  candle,  oil  lamp,  or  alcohol  lamp 
must  be  used,  else  the  sulphur  in  the  burning  gas  will 
vitiate  the  results. 


RECAPITULATION   OF  THE   EXAMINATION  ON 
CHARCOAL. 

The  substance  is  heated  on  charcoal. 

DEFLAGRATION. — Nitrates  and  chlorates  (14). 

FUSION. — Potassium  and  sodium  salts  (15). 

WHITE  INFUSIBLE  RESIDUE. — Compounds  of  metals  of  Groups 
II.  and  III.,  zinc  salts  and  silicic  acid  (10). 

DARK-COLORED  RESIDUE. — Compounds  of  metals  of  Groups 
IV.  and  V.,  except  zinc  and  mercury  (17). 

COMPLETE  VOLATILIZATION. — Ammonic  and  mercuric  com- 
pounds, and  compounds  of  arsenic  and  antimony,  which 
contain  no  other  metal  (18). 

BRIGHT  METALLIC  GLOBULES. — Silver  and  gold  compounds, 
and  the  oxides  of  lead  and  bismuth  (19). 

The  substance  is  mixed  with  sodic  carbonate  and  heated  on 
charcoal. 

METALLIC  GLOBULES  WITHOUT  INCRUSTATION. — Gold,  silver, 
and  copper  (22)  ;  tin  (24). 

METALLIC  GLOBULES  AND  INCRUSTATION. — Lead  (25) ;  bis- 
muth (26)  ;  antimony  (28). 


9o  PART  III. 

INCRUSTATION  WITHOUT  GLOBULES. — Zinc  (23)  ;  Cadmium 

(27).     GARLIC  ODOR.— Arsenic  (29). 
FORMATION  OF  SODIC  SULPHIDE. — All  compounds  containing 

sulphur  (30). 


PRELIMINARY   TESTS   WITH  METALLIC  BODIES. 

EXAMINATION  IN  A  CLOSED  TUBE. 
See  page  26. 

Mercury.      Amalgams   containing  mercury  give  a  subli- 

(31)  mate  of  METALLIC  MERCURY  when  they  are  heated. 
At  first  a  gray  film  forms  in  the  upper  part  of  the  tube, 
and,  when  the  amount  of  mercury  is  considerable,  fine 
globules  of   metallic  mercury  are  formed,  which  ag- 
glomerate and  become  more  distinctly  visible  when 
they  are  rubbed  with  a  copper  wire. 

Arsenic*     Some  metallic  compounds,  containing  ARSENIC, 

(32)  give  a  metallic  mirror  or  ring  in  the  upper  part  of  the 
tube  when  they  are  heated. 


EXAMINATION  ON  CHARCOAL. 

Heat  a  piece,  one-fourth  as  large  as  a  pea,  of  the 
(33)  metallic  substance  in  a  cavity  on  a  piece  of  charcoal. 
See  Part  I.,  page  25.  The  phenomena  to  be  ob- 
served are  the  formation  of  the  incrustations  de- 
scribed (pages  87  and  88),  the  smell  of  ARSENIC,  and 
the  vapors  of  MERCURY  and  ANTIMONY. 

COPPER  colors  the  blowpipe  flame  green,  or  in  the 
presence  of  chlorine  blue. 


SULPHURIC  ACID   TEST.  9! 

PRELIMINARY  TESTS  WITH  NON-METALLIC 
SOLIDS  (continued}. 

TESTS  WITH  THE  BORAX  BEAD. 

If  the  substance  to  be  tested  appears  to  be  the  oxide,  or  an 
oxygen-salt  of  a  higher  metal  (see  17),  dissolve  some  of  it  in 
the  borax  bead. 
Cobalt  colors  the  bead  blue  in  the  oxidizing  and  in  the  re- 

(34)  ducing  flame. 

Copper  colors  the  bead  green  when  hot,  and  blue  when  cold, 

(35)  in  the  oxidizing  flame.     It  colors  the  bead  red,  when 
cold,  in  the  reducing  flame. 

Chromium  colors  the  bead  green  in  both  flames. 

(36) 
Iron*  colors  the  bead  brownish  red  when  hot,  and  yellow  when 

(37)  cold,  in  the  oxidizing  flame. 

Nickel  colors  the  bead  violet  when  hot,  and/0/<?  brown  when 

(38)  cold,  in  the  oxidizing  flame.     The  color  disappears  in 
a  good  reducing  flame.     (See  page  44.) 

IVLaifigd'nese  colors  the  bead  amethyst  in  the  oxidizing  flame. 

(39)  The  color  disappears  in  a  good  reducing  flame. 

(40)  THE  OXYGEN    COMBINATIONS   OF  THE  REMAINING 
METALS  color  the  bead  very  slightly  or  not  at  all. 

CONCENTRATED  SULPHURIC  ACID  TEST. 

CONCENTRATED  SULPHURIC  ACID,  with  the  aid  of  heat,  sets 
free  other  acids  from  most  of  their  combinations  with  metals, 
and  frequently  in  such  a  form  that  they  can  be  recognized  by 
simple  tests.  This  reaction  is  not,  of  course,  a  method  of 
separation  ;  one  acid  may  obscure  the  test  for  another,  and 
the  possible  cases  are  so  complicated  that  it  would  be  useless 
to  attempt  to  describe  them  all ;  therefore,  if  the  result  of  the 


92 


PART  III. 


sulphuric  acid  test  appears  doubtful,  it  is  best  to  reserve  judg- 
ment of  its  value  until  after  the  tests  for  acids  in  solution  have 
been  applied. 

The  reactions  of  acids  or  their  salts,  when  a  small  quantity 
of  a  solid  substance  or  of  a  very  concentrated  solution  is  added 
to  a  few  cubic  centimetres  of  strong  sulphuric  acid  in  a  test- 
tube  are  described  below.  Heat  should  be  applied  after  the 
reaction,  which  takes  place  at  the  ordinary  temperature,  has 
been  observed. 
Carbonic  Acid.  Effervescence.  Carbonic  dioxide  gas 

(41)  renders  turbid  a  drop  of  lime-water  held  on  the  end 
of  a  glass  rod  in  the  test-tube  (see  page  70,  foot-note). 
Carbonic  acid  is  also  detected  by  the  chlorhydric  acid 
test,  and  in  many  cases  that  test  is  preferable  to  the 
one  with  sulphuric  acid,  since  oxalic  acid  does  not 
give  the  same  reaction  with  chlorhydric  acid.     See 

(73). 

Oxalic  Acid.     When  a  dry  compound  of  oxalic  acid  is 

(42)  added    to    strong    sulphuric    acid,    and   the   mixture 
is  heated,  carbonous  oxide  and  carbonic   oxide  are 
evolved. 

The  sulphate  of  calcium  test  for  oxalic  acid  (125) 
is  more  accurate  than  that  with  sulphuric  acid. 
Cyarihydric,     Ferro-    and    Ferricyanhydric 

(43)  Acids.     Compounds    of  these   acids   evolve,  when 
perfectly  dry,  carbonic  oxide  with  effervescence  when 
they  are  heated  with  strong  sulphuric  acid.     Usually, 
however,  a  faint  odor  of  cyanhydric  acid  can  be  de- 
tected.    The  special  tests   are    more  valuable.     See 
(75),  and  (133),  (134),  and  (135). 

Fluorhydric  Acid.     When  a  solid  substance  or  a  con- 

(44)  centrated  solution  containing  fluorine  is  heated  with 
strong  sulphuric  acid,  fluorhydric  acid  is  evolved,  which 
etches  glass  (see  page  69)  ;  when  silicic  acid  or  a  sili- 
cate is  present,  fumes  of  fluoride  of  silicon,  which  give 


SULPHURIC  ACID  TEST.  93 

a  precipitate  of  silica  in  a  drop  of  water  held  over 
them  on  the  end  of  a  glass  rod,  are  evolved. 

If  fluorhydric  acid  is  discovered,  the  substance 
which  is  to  be  used  for  page  104,  III.,  and  the  follow- 
ing tests,  must  be  heated  with  sulphuric  acid  in  a  pla- 
tinum vessel  until  the  fluorhydric  acid  is  driven  off 
completely. 

A  separate  portion  can  be  used  for  testing  for  sul- 
phuric acid.  See  (30). 

Sulphydric  Acid.  Compounds  containing  this  acid 
(4#)  evolve  it  (often  with  effervescence)  when  strong  sul- 
phuric acid  is  added  to  them.  Sulphydric  acid  may  be 
recognized  by  its  smell  and  by  its  property  of  black- 
ening paper  dipped  in  plumbic  acetate  solution.  See 
(74)  and  (129). 

Sulphurous   Acid.     Compounds   containing    this    acid 

(40)  evolve  sulphurous  oxide,  SO2,  with  effervescence  when 

strong  sulphuric  acid  is  added  to  them.     Sulphurous 

oxide,  when  free  from  sulphydric  acid,  and  from  some 

others,  can  be  recognized  by  its  smell.     See  (7#). 

Chloric  and  Hypochlorous  Acids.    (See  page  80.) 

(47)  Compounds  containing  these  acids  evolve  a. yellow  gas 
on  the  addition  of  strong  sulphuric  acid,  even  when 
the  mixture  is  not  heated.    The  gas  can  best  be  recog- 
nized by  its  color,  its  odor,  and  its  strong  bleaching 
action  on  a  drop  of  indigo  solution  held  in  the  tube 
on  the  end  of  a  glass  rod.    This  gas  precipitates  nitrate 
of  silver. 

Chlorhydric  Acid.     Compounds  containing  chlorhydric 

(48)  acid  evolve  the    gas,  HC1,  frequently  with  efferves- 
cence, on  the  addition  of   sulphuric  acid.     The  gas 
does  not  bleach  indigo  solution,  but  precipitates  ni- 
trate of  silver  held  on  the  end  of  a  glass  rod  in  the 
tube.     Chlorhydric  acid,  in  the  presence  of  an  oxidiz- 
ing agent,  evolves  chlorine  under  the  same  circum- 


94  PART  in. 

stances.  The  gas  produces  the  same  reactions  as  are 
produced  by  the  gas  evolved  by  chloric  and  hypo- 
chlorous  acids,  but  it  can  be  distinguished  from  them 
by  its  smell  and  by  its  color,  which  is  a  less  intense 
yellow.  For  special  test,  see  (130). 
Hromhydric  Acid.  Bromides  are  decomposed  by  strong 

(49)  sulphuric   acid  with    evolution   of  bromhydric   acid, 
which,  if  the  sulphuric  acid  is  concentrated  and  in  ex- 
cess, is  partly  decomposed,  with  separation  of  bromine 
and  formation  of  sulphurous  oxide.  (See  also  page  73.) 

lodhydric  Acid.     All    the  iodides   are  decomposed  by 

(50)  strong    sulphuric    acid  on   the    application   of    heat. 
Iodine  is  set  free,  which  escapes  in  violet  vapors  and 
imparts  a  blue  color  to  paper  moistened  with  starch. 
(See  also  page  73.) 

Nitric  Acid.     Compounds  containing  nitric  acid  in  con- 

(51)  siderable  quantity  produce  reddish  fumes  when  heated 
with  sulphuric  acid  in  the  presence  of  copper  turnings 
or  of  any  other  reducing  agent.     The  following  test 
for  nitric  acid  is  more  delicate  :  Mix  a  little  of  the 
powdered  substance  or  solution  with  strong  sulphuric 
acid,  and  pour  cautiously  upon  the  acid  a  solution  of 
ferrous  sulphate.     A  brown  or  red  color  at  the  line  of 
separation  of  the  two  solutions  indicates  the  presence 
of  NITRIC  ACID,  HNO8.     See  (136)  (a)  for  this  test 
in  the  presence  of  ferro-  or  ferricyanhydric  acid. 

Acetic  Acid  gives  with  sulphuric  acid  an  odor  of  vine- 
gar.  The  following  test  is  more  delicate  :  Add  an 
equal  volume  of  alcohol  to  strong  sulphuric  acid,  and 
then  add  the  solid  substance  or  concentrated  solution 
supposed  to  contain  ACETIC  ACID,  C2H4O2.  If  this 
acid  is  present,  the  odor  of  acetic  ether  will  be  per- 
ceptible on  heating  the  mixture.  It  is  well  to  add 
pure  acetic  acid  at  the  same  time  to  a  mixture  of  sul- 
phuric acid  and  alcohol,  in  order  to  compare  the  odor 


SOL  UTION  OF  NON-ME TALLIC  BODIES.  95 

produced  with  that  observed  in  the  test.  If  other 
gases  render  the  odor  of  acetic  ether  difficult  to  per- 
ceive, the  precautions  described  on  page  81  must  be 
observed. 

RECAPITULATION  OF  SULPHURIC  ACID  TEST. 

A  colorless  gas  is  given  off. 

THE  GAS  is  WITHOUT   ODOR. — Carbonic  acid  (4:1)  ;  oxalic 

acid  (42). 
A  PUNGENT  SUFFOCATING  ODOR. — Fluorhydric  acid  (44)  ; 

chlorhydric  acid  (48)  ;  bromhydric  acid  (40). 
AN  ODOR  OF  BITTER  ALMONDS. — Cyanhydric,  ferrocyanhydric, 

and  ferricyanhydric  acids  (43). 
AN  ODOR  OF  ROTTEN  EGGS. — Sulphydric  acid  (43). 
AN  ODOR  OF  BURNING  SULPHUR.— Sulphurous  acid  (4:6). 
AN  ODOR  OF  VINEGAR. — Acetic  acid  (52). 

A  colored  gas  is  given  off. 
*  The  gas  has  also  a  peculiar  suffocating  odor. 

YELLOW  GAS. — Chloric  and  hypochlorous  acids  (4:7). 
FAINT  YELLOW  GAS. — Chlorine  (48). 
RED  FUMES. — Nitric  acid  with  copper  (SI). 
VIOLET  FUMES. — lodhydric  acid  (50). 

METHODS    OF    DISSOLVING    NON-METALLIC 
BODIES. 

Substances  to  be  tested  may  be  divided  into  three  classes. 

1st  Class  :     Bodies  soluble  in  water. 

To  ascertain  whether  a  body  is  entirely  soluble, 
take  a  few  grains  of  the  substance,  which,  when  it  dis- 
solves with  difficulty,  must  be  in  as  finely  divided  a 


96  PART  III* 

condition  as  possible,  and  digest  them  with  a  con- 
siderable quantity  of  water  in  a  test-tube.  If  the  sub- 
stance does  not  dissolve  completely,  boil  it  for  a  few 
minutes  with  water.  If  it  still  does  not  dissolve  com- 
pletely, filter  and  evaporate  a  few  drops  of  the  filtrate 
on  platinum  foil,  in  order  to  see  whether  anything  has 
dissolved. 

If  the  substance  is  partly  soluble  in  water,  treat  a 
considerable  quantity  in  the  manner  directed  above, 
repeat  the  boiling  with  water  once  or  twice,  wash 
thoroughly  with  water,  and  filter,  taking  care  that  as 
little  as  possible  of  the  solid  substance  goes  upon  the  fil- 
ter, and  use  the  insoluble  portion  for  the  following  test. 

A  separate  analysis  should  usually  be  made  of  the 
part  of  a  substance  which  is  soluble,  and  of  that  which 
is  insoluble  in  water. 

If  the  substance  is  wholly  insoluble  in  water,  it  may 
be  used  immediately  for  the  following  test.  It  should 
usually  be  finely  pulverized. 

2d  Class  :  Bodies  insoluble  in  water,  but  soluble  in  chlor- 
(54:)  hydric  acid  or  in  nitric  acid  or  in  aqua  regia.  Digest 
a  few  grains,  or  a  very  small  quantity  of  the  finely 
powdered  substance  with  dilute  chlorhydric  acid.  If 
it  does  not  entirely  dissolve,  boil  it  with  the  acid.  If 
it  still  does  not  entirely  dissolve,  pour  off  the  liquid, 
and  boil  the  insoluble  substance  with  strong  chlorhy- 
dric acid.  If  the  substance  is  not  wholly  soluble  in 
chlorhydric  acid,  repeat  the  trial  with  nitric  acid  in 
the  same  way,  using  a  fresh  portion  of  the  substance. 
If  the  substance  is  still  insoluble,  use  a  mixture  of  both 
.acids  (aqua  regia). 

If  sulphur,  which  may  be  recognized  by  its  color  and 
its  low  specific  gravity,  or  silicic  acid,  which  may  be  re- 
cognized by  its  peculiar  gelatinous  aspect,  separate  out  on 
the  addition  of  chlorhydric  or  of  nitric  acids,  the  sub- 


SOLUTION  OF  NON-METALLIC  BODIES.  97 

stance  must  be  considered  as  soluble,  and  after  filtration 
(in  case  of  silicic  acid,  see  64)  the  solution  must  be  tested 
according  to  page  104  (///.),  or  pag'e  107  (VI r.),  accord- 
ing as  nitric  or  chlorhydric  acid  has  been  the  solvent. 

If  the  substance  is  only  partially  soluble  after  treat- 
ment as  above,  the  insoluble  portion  must  be  washed 
carefully  and  separated  by  filtration  from  the  soluble 
portion,  and  it  must  be  subjected  to  the  tests  for  the 
3d  Class.  See  page  137. 

In  all  cases  before  proceeding  to  test  the  solution 
obtained  by  treatment  with  acids,  it  must  be  so  pre- 
pared that  it  will  be  moderately  acid  with  chlorhydric 
acid,  and  will  not  contain  any  free  nitric  acid,  or  great 
excess  of  strong  acid  of  any  kind.  It  the  solution  was 
made  by  the  use  of  dilute  chlorhydric  acid  it  is  already 
in  this  condition.  If  nitric  acid  was  used  for  the  solu- 
tion the  excess  must  be  replaced  by  chlorhydric  acid. 
In  this  case  add  a  quantity  of  the  latter  about  half 
as  great  as  that  of  the  nitric  acid  used  ;  evaporate  the 
liquid  almost  to  dryness  in  the  hood,  and  add  from  25 
to  50  c.c.  of  water  to  the  residue.  If  the  substance 
was  dissolved  in  concentrated  chlorhydric  acid,  the 
solution  should  be  evaporated  down  in  the  same 
manner  to  expel  the  excess  of  that  acid,  and  the  resi- 
due diluted  considerably  with  water.  If  the  addition 
of  chlorhydric  acid  should  cause  a  precipitate  in  the 
clear  acid  solution,  or  if  the  residue,  after  evaporation, 
does  not  dissolve  entirely  in  water,  it  can  only  be  occa- 
sioned by  the  presence  of  argentic,  plumbic,  or  mercu- 
rous  chloride.  In  this  case,  after  the  addition  of  water, 
filter  off  the  insoluble  powder  and  examine  for  those 
metals  as  directed  (OS),  (67),  (69).  Use  the  fil- 
trate for  the  sulphydric  acid  test  (76). 

3(1  Class  ;     Bodies  insoluble  in  water,  and  in  chlorhydric, 
nitric,  and  nitro-chlorhydric  acids  (aqua  regia). 
7 


98  PART  III. 

The  tests  to  be  applied  to  bodies  of  this  class  follow 
those  in  the  scheme  for  testing  bodies  in  solution.  See 
Page  137. 

METHODS  OF  DISSOLVING  METALLIC  BODIES. 
Metals  are  divided  into  three  classes. 

1st  Class  :     Metals  which  are  not  attacked  by  NITRIC  ACID. 
(50)       If  the  metal  does  not  appear  to  be  entirely  soluble 
in  dilute  nitric  acid,  even  after  boiling,  use  strong  ni- 
tric acid. 

GOLD  is  insoluble,  also  alloys  containing  a  very  large 
proportion  of  gold. 

2d  Class  :     Metals  which  are  attacked  by  nitric  acid,  and 
(«57)  converted  into  oxides  (white  powder],  which  are  inso- 
luble in  the  acid. 

TIN  and  ANTIMONY  belong  to  Class  II. 

3d  Class  :  Metals  which  dissolve  entirely  in  nitric  acid. 
(58)  All  the  remaining  metals  dissolve  entirely  in  nitric 
acid.  With  metals  of  this  class  the  solution  should  be 
effected  by  heating  with  nitric  acid  in  an  evaporating 
dish  until  no  more  red  fumes  are  given  off.  Chlorhy- 
dric  acid  should  then  be  added,  the  greater  part  of  the 
acid  evaporated,  water  added,  and  the  solution  should 
be  tested  according  to  (page  104,  III.).  If  the  metal 
remains  partly  insoluble,  a  small  portion  of  it  should 
be  carefully  tested  by  boiling  with  strong  nitric  acid, 
to  see  whether  it  will  not  dissolve  by  using  a  sufficient 
quantity  of  the  acid. 

An  alloy  may  contain  metals  of  each  class,  therefore, 
after  the  treatment  with  nitric  acid  described  in  the 
last  paragraph,  if  an  insoluble  residue  is  found  it  should 
be  dissolved  in  aqua  regia,  and  the  solution  boiled 
until  the  smell  of  chlorine  ceases  to  be  given  off. 
Then  if  the  residue  had  a  metallic  appearance  the  solu- 


SOLUTION  OF  METALLIC  BODIES.  99 

tion  must  be  tested  for  gold.  See  (87).  If  it  was  a 
white  powder  it  must  be  tested  for  tin  and  antimony. 
See  (86)  and  (85). 

SOLUTION  IN  CHLORHYDRIC  ACID. 

(00)  Sometimes  a  test  shows  that  the  metallic  body  can 
be  readily  dissolved  in  chlorjiydric  acid  (see  ZINC, 
page  39,  "and  IRON,  page  41),  and  in  this  case  the  chlor- 
hydric  acid  solution  should  be  preferred,  and  it  should 
be  tested  according  to  page  107,  VI. 


TESTS    FOR    METALS    AND 
ACIDS. 


BODIES  IN  SOLUTION. 
CLASSES  I.  AND  II.     (See  page  95.) 
Sttbstances  dissolved  in  water  or  in  acids. 

(The  tests  I.  and  II.,  which  follow,  are  only  to  be  used  for 
substances  dissolved  in  water,  or  where  the  solvent  is 
unknown.) 

I.— REACTION  WITH  TEST-PAPER. 

(01)  Observe  the  reaction  with  litmus  paper. 

(02)  If  the  reaction  is  acid,  add  to  a  small  portion  of  the 
solution  sodic  carbonate,  drop  by  drop,  until  the  effer- 
vescence ceases,  and  then  heat  to  boiling. 

If  metals  of  the  2d,  3d,.  4th,  and  5th  groups  are 
present,  a  precipitate  will  be  formed,  except  in  a  few 
special  cases. 

Some  idea  of  the  amount  of  free  acid  present  in  the 
solution  can  be  formed  by  observing  the  amount  of 
sodic  carbonate  required  to  neutralize  it. 

II.— EVAPORATION. 

(03)  Test  one  or  two  drops  of  the  solution  by  evaporation 
on  platinum  foil  (or  on  a  bit  of  glass  or  porcelain,  if 
there   is  danger  of  injury  to  the  platinum),  to  see 
whether  it  leaves  a  residue. 

100 


PRELIMINARY  EXAMINATION. 

If  this  test  shows  the  presence  of  solid  matter  in  the 
solution,  a  portion  of  the  latter  may  be  evaporated  to 
dryness  in  a  small  porcelain  dish,  and  some  of  the  tests 
for  solids,  particularly  the  EXAMINATION  ON  CHARCOAL 
(page  85),  and  the  CONCENTRATED  SULPHURIC  ACID 
TEST  (page  91),  may  be  applied  to  the  dry  substance 
thus  obtained.  It  is  best  only  to  use  the  evaporation 
and  tests  applied  to  the  dry  substance,  to  settle  any 
doubts  that  remain  in  regard  to  the  constitution  of  the 
solution,  after  the  tests  usually  applied  to  solutions 
have  been  performed.  (See  page  104,  etc.) 

When  a  solution  of  unknown  origin  is  presented  for 
analysis  it  should  always  be  heated,  in  order  to  see 
whether  a  gas  is  given  off  that  can  be  recognized  by 
the  tests  (page  92,  and  page  93). 

Silicic  Acid  is  usually  not  present,  except  in  alkaline  solu- 
tions ;  it  occurs,  however,  in  small  quantities  in  spring 
and  river  waters,  and  it  may  also  exist  in  acid  solu- 
tions. 

Unless  silicic  acid  is  known  to  "be  absent  it  should 
be  tested  for,  and  removed  before  making  the  remain- 
ing tests.  To  this  end  render  the  solution  acid  with 
chlorhydric  acid,  and  evaporate  carefully  to  dryness 
in  a  small  porcelain  dish.  Care  must  be  taken  not  to 
heat  the  dish  over  the  lamp  after  its  contents  have  be- 
come dry.  Treat  the  dry  substance  with  a  few  drops 
of  acid,  and  then  boil  with  water.  If  there  is  an  inso- 
luble residue,  it  consists  of  silica,  SiO2.  (See  page  71.) 

Subject  the  chlorhydric  acid  solution  to  the  remain- 
ing tests,  beginning  with  page  107,  VI.  ;  and  if  there 
is  reason  to  suspect  that  the  precipitate,  left  after  the 
treatment  with  an  acid,  contains  other  substances  be- 
sides silica,  examine  it  according  to  page  105,  IV. 


PART  III. 


TESTS  FOR  METALS. 


THE  following  scheme  of  testing  for  metals  is  founded  upon 
the  successive  precipitation  of  a  number  of  groups,  which  in- 
clude all  the  metals  as  far  as  Group  II.  After  the  metals  of 
the  higher  groups  have  been  removed  by  precipitation,  or  have 
been  found  to  be  absent,  those  of  Group  II.  are  precipitated 
successively.  Sodium'  and  potassium  are  detected  by  the 
colors  of  their  flames  in  a  solution  which  has  been  freed  from 
all  the  higher  metals  except  magnesium.  Ammonia  can  be 
detected  in  a  solution  without  having  reference  to  its  other 
constituents. 

After  the  separation  into  groups  has  taken  place  by  precipi- 
tation with  the  general  reagents,  each  precipitate,  which  may 
contain  one  or  all  the  metals  belonging  to  its  group,  is  usually 
dissolved,  and  the  further  analysis  is  performed  by  testing  in 
the  several  solutions  for  all  the  metals  which  they  may  contain. 
These  special  tests  sometimes  require  the  separation  of  the  dif- 
ferent metals,  one  after  another,  in  a  particular  order,  while 
sometimes  a  test  for  a  metal  may  be  applied  to  the  solution 
without  regard  to  the  presence  of  other  metals.  In  all  cases 
the  conditions  requisite  for  applying  the  tests  will  be  described. 

The  general  tests  must  be  applied  in  the  following  order : 
ist,  Chlorhydric  acid  to  effect  the  precipitation  of  silver  and 
mercurous  compounds,  and  of  lead  if  it  is  present  in  large 
quantity.  2d,  Sulphydric  acid  in  an  acid  solution  to  precipi- 
tate small  quantities  of  lead  and  the  metals  of  Group  V.,  Sec- 
tion II.,  and  of  Group  VI.  (From  this  precipitate  the  metals 
of  Group  VI.  are  separated  by  dissolving  them  in  ammonic  sul- 


TESTS  FOR  METALS.  IO3 

phide.)  36,  Ammonic  hydrate  until  the  reaction  becomes 
alkaline,  ammonic  chloride  and  ammonic  sulphide  to  precipi- 
tate the  metals  of  Groups  III.  and  IV. 

Whenever  a  single  metal  or  a  group  of  metals  is  precipitated 
some  of  the  liquid  containing  the  precipitate  must  be  poured 
on  a  filter,  and  the  first  drops  of  the  solution  which  runx 
through  must  be  tested  with  some  of  the  reagent  which  was 
used  to  produce  the  precipitation,  in  order  to  ascertain  whether 
it  has  been  completely  effected.  Should  a  fresh  precipitate 
make  its  appearance,  everything  must  be  poured  back  from 
the  filter  into  the  test-tube  or  flask  and  more  of  the  reagent 
must  be  added.  This  operation  must  be  repeated  until  no 
precipitation  is  produced  by  the  same  reagent  in  the  liquid 
which  runs  through  the  filter.  After  a  little  practice  it  is  easy 
to  estimate  how  much  of  a  reagent  is  required  to  effect  a  com- 
plete precipitation.  A  surplus  over  this  quantity  is  called  an 
excess  of  the  reagent. 

When  it  has  been  found  that  a  slight  excess  of  the  reagent 
has  been  added,  the  whole  of  the  liquid  and  the  precipitate 
together  must  be  poured  upon  a  fresh  filter  and  the  liquid 
must  be  allowed  to  drain  off.  The  liquid  is  to  be  tested  for 
metals  of  the  succeeding  groups,  and  the  precipitate  must 
usually  be  completely  freed  from  it ;  otherwise  the  separation 
has  no  value.  To  this  end  water  must  be  blown  on  the  preci- 
pitate from  the  wash-bottle  and  allowed  to  drain  off.  Care 
must  be  taken  not  to  let  the  water  overflow  the  edge  of  the 
filter.  The  washing  must  be  continued  until  the  water  which 
flows  through  the  filter  is  proved,  either  by  evaporation  on 
platinum  foil  or  by  the  application  of  tests  for  the  succeeding 
groups,  not  to  contain  any  metal  in  solution.  The  last  portions 
of  the  wash-water  which  pass  through  the  filter  may  be 
thrown  away,  as  they  contain  very  little  of  the  substances  to 
be  tested  for. 

The  value  of  analyses  depends  upon  the  care  with  which 
the  separation  of  precipitates  from  the  liquid  in  which  they 


104 


PART  III. 


have  been   formed  is  executed.     For   special   directions  for 
filtering,  see  pages  23  and  24. 

When  the  reaction  of  a  liquid  in  a  test-tube  is  to  be  tested, 
always  close  the  tube  and  shake  it  thoroughly  before  dipping 
the  test-paper  in. 


METALS  OF  GROUPS   VI.   AND   V. 

III.— CHLORHYDRIC  ACID  TEST. 

Metals  in  acid  solution. 

GROUP  V.,  SECTION  I. 

In  case  the  solution  is  known  to  contain  chlorhydric  acid,  pass  to 
page  107  (VI.). 

If  the  solution  has  an  alkaline  reaction,  pass  to  page  105  (IV.). 

(63)  Add  to  a  very  small  quantity  of  the  solution  to  be 
tested  a  few  drops  of  dilute  chlorhydric  acid.  If  a 
precipitate  forms,  continue  to  add  chlorhydric  acid, 
drop  by  drop,  as  long  as  it  seems  to  increase  in  quan- 
tity, then  add  a  quantity  of  chlorhydric  acid  about 
equal  to  that  already  added. 

If  no  precipitate  is  formed,  or  if  it  is  dissolved  on 
further  addition  of  dilute   chlorhydric  acid,    LEAD  IN 

LARGE    QUANTITY,    SILVER,     AND    MERCUROUS    SALTS 

ARE  ABSENT.     Pass  to  page  107  (VI.). 
(06)       If  a  permanent  precipitate  is  formed,  it  may  consist 

Of  PLUMBIC,    ARGENTIC,  AND    MERCUROUS    CHLORIDES. 

Most  of  the  succeeding  tests  for  the  detection  of  bases 
must  be  performed  in  a  solution  freed  from  these 
metals.  Therefore  if  a  precipitate  has  been  observed, 
treat  a  considerable  quantity  of  the  solution  in  the 
same  way  that  the  small  portion  was  treated,  and  after 
an  excess  of  chlorhydric  acid  has  been  added  shake 


CHLORHYDRIC  ACID  TEST. 


105 


the  liquid  for  a  minute  or  two.  The  precipitate  will 
then  settle  in  a  short  time,  leaving  the  solution  nearly 
clear.  The  solution  should  be  decanted  through  a 
filter,  and  the  precipitate  washed  twice  by  decantation 
through  the  filter  with  water  acidulated  with  chlor- 
hydric  acid. 

Test  the  filtrate  according  to  page  107  (VI.). 

Lead.      Add  a  small  quantity  of  water  to  the  precipitate^ 

(67)  and  boil,  then  let  the  precipitate  settle,  and  decant  the 
clear  liquid  through  the  same  filter  which  was  used  in 
(66)  into  another  vessel.     Add  to  the  filtrate  an  equal 
bulk  of  alcohol  and  a  small  quantity  of  dilute  sul- 
phuric acid.     If  a  white  precipitate  forms  it  consists 

Of   SULPHATE  OF  LEAD,  PbSO4. 

If  lead  is  found,  the  precipitate  must  be  washed  as 
before,  with  boiling  water  by  decantation,  until  the 
filtrate  gives  no  black  precipitate  with  ammonic  sul- 
phide. If  no  precipitate  remains  after  the  washing,  NO 
ARGENTIC  OR  MERCUROUS  SALTS  are  present.  Pass  to 
page  107  (VI.). 
Silver.  If  a  precipitate  remains  after  washing  with  boiling 

(68)  water,  add  to  it  ammonia,  pour  the  solution  through 
a  filter,  and  acidify  with  nitric  acid.     If  a  precipitate 
forms,  it  consists  of  ARGENTIC  CHLORIDE,  AgCl.. 

ftfercurous  Salts.     If  a  precipitate  remains  after  ammo- 

(69)  nia  has  been  added,  it  will  have  a  gray  or  black  color. 
The  precipitate  consists  of  A  MERCUROUS  COMPOUND 

OF  AMMONIA. 

IV— CHLORHYDRIC  ACID  TEST. 

Metals  in  alkaline  solutions. 

(70)  Add  chlorhydric  acid  until  the  reaction  becomes 
distinctly   acid,  and  if  a  precipitate   forms,  wash  it 
thoroughly  with  cold  water  upon  a  filter  until  the  fil- 


I06  PART  III. 

trate  is  no  longer  acid  to  test-paper.  Observe  whether 
sulphydric  acid  is  given  off. 

If  a  precipitate  is  formed  with  chlorhydric  acid,  filter 
and  test  the  filtrate  according  &>page  107  (VI.). 
(71)       If  no  precipitate  is  formed,  pass  to  page  107  (VI.). 

(a)  If  the  precipitate  is  white,  it  may  consist  of : 
Plumbic  Chloride.     Boil  a  little  of  it  with  water,  and 

test  one  portion  of  the  solution  for  LEAD  according  to 
(07))  and  test  another  portion  for  CHLORINE  by  add- 
ing nitric  acid  and  argentic  nitrate. 

Plumbic  Sulphate.     Test  according  to  (140). 

Argentic  Chloride.    Test  according  to  (141). 

(b)  If  the  precipitate  is  colored  it  may  contain 

The  Sulphides  of  Arsenic,  Antimony,  and  Tin. 
Test  according  to  page  109  (VIII.). 

Sulphur  may  be  precipitated,  accompanied  by  a  disengage- 
ment of  sulphydric  acid.  The  precipitated  sulphur 
can  be  recognized  by  its  appearance,  and  its  insolu- 
bility in  aqua  regia. 

V.— CHLORHYDRIC  ACID  TEST  FOR  ACIDS. 
See  also  Silicic  Acid  (64). 

If  a  solution  is  acid,  many  of  the  following  tests, 
particularly  (73),  (74),  and  (7S),  can  be  applied 
by  simply  heating  the  solution. 

Carbonic  Acid  is  only  present  in  alkaline  solutions.  CO2 
is  evolved  with  effervescence,  when  an  acid  is  added, 
until  the  solution  has  an  acid  reaction.  Hold  a  drop 
of  lime-water  on  the  end  of  a  glass  rod  in  the  tube  ;  if 
CARBONIC  DIOXIDE,  CO2,  is  present,  a  white  precipitate 
forms.  (See  page  70,  foot-note.)  The  same  test  can 
be  applied  to  solid  carbonates.  See  also  the  sulphuric 
acid  test  (41). 

The  following  acids  need  only  be  looked  for  when  an 


S ULPH YDRIC  A  CID  TEST.  ! o 7 

odor  can  be  perceived  after  heating  the  solution,  or  after 
adding  chlorhydric  acid  and  heating. 
Cyanhydric    Add,    in   its   soluble   combinations  with 

(73)  most  metals,  is  set  free  by  chlorhydric  acid.     It  can 
be  recognized  by  its  smell.     See  also  the  prussian  blue 
test  (133), 

Sulphydric   Acid    is    evolved   from    alkaline    solutions 

(74)  (often  with  effervescence),  on  the  addition  of  chlorhy- 
dric acid,  when  the  solution  is  heated.     It  can  be  re- 
cognized by   its  smell  and  by  the  lead  paper  test. 
(See  the  sulphuric  acid  test  (4#)   and  the  argentic 
nitrate  test  (129). 

SulpJlllTOUS  Oxide  is  evolved  from  alkaline,  neutral,  or 
(7«>)  slightly  acid  solutions  of  sulphites,  on  the  addition  of 
chlorhydric  acid.  Mix  a  little  potassic  ferricyanide 
and  ferric  chloride,  and  hold  a  drop  of  the  mixture  on 
the  end  of  a  glass  rod  in  the  tube  after  chlorhydric 
acid  has  been  added  and  the  tube  heated.  If  a  blue 
color  appears,  SULPHUROUS  ACID  is  present.  If  sul- 
phydric  acid  is  present  add  sufficient  plumbic  acetate 
to  precipitate  it  before  performing  the  test.  See  the 
sulphuric  acid  test  (40). 

VI.— SULPHYDRIC  ACID  TEST. 

Metals  in  acid  solutions. 

(76)  Add  sulphydric  acid  solution  to  a  small  quantity  of 
the  solution  to  be  tested,  and  warm  gently.  In  case 
metals  of  Group  VI.  are  to  be  tested  for,  it  is  better 
to  pass  sulphydric  acid  gas  into  the  dilute  solution, 
made  acid  with  chlorhydric  acid.  The  total  precipi- 
tation of  the  metals  of  Group  VI.  is  frequently  only 
effected  after  one  or  two  days. 

If  no  precipitate  forms,  no  metals  of  Grozips  V.  and 
VI.  are  present.     Pass  to  page  115  (X.). 


I08  PART  III. 

(77)  If  a  precipitate  forms  observe  the  color.    It  may  con- 
sist Of  the  SULPHIDES  OF  LEAD,  PbS  J    BISMUTH,  Bi2S8  J 

COPPER,  CuS ;  MERCURY,  HgS  ;  and  GOLD,  Au2S3, 
when  it  is  black  ;  ARSENIC,  As2S3 ;  TIN  (BISULPHIDE), 
SnS2 ;  CADMIUM,  Cd$>,  yellow  j  TIN  (MONOSULPHIDE), 
SnS,  brown  ;  ANTIMONY,  Sb2S3  or  Sb2S5,  orange.  The 
presence  of  a  black  sulphide  hides  the  color  of  the 
other  sulphides,  so  that  all  may  be  present  when  the 
precipitate  is  black. 

If  only  a  light,  fine,  white  precipitate,  which  is  not  de- 
stroyed by  acids,  is  formed,  it  consists  of  sulphur,  and  is 
frequently  due  to  the  presence  of  a  ferric  salt  or  a  chro- 
mate  in  the  solution.  In  case  only  sulphur  is  precipi- 
tated, pass  to  page  115  (X.). 

(78)  If  a  precipitate  forms  in  a  small  portion  of  the  solu- 
tion, a  sufficient  quantity  for  use  in  all  the  succeeding 
tests  for  metals  must  be  treated  with  sulphydric  acid 
until  the  metals  of  Groups  V.  and  VI.  are  completely 
precipitated  as  sulphides  ;    and  the  precipitate  thus 
obtained   must  be  washed  on   a  filter  quickly,  with 
warm  water  containing  sulphydric  acid,  until  the  ad- 
dition of  ammonic  hydrate  to  the  filtrate  ceases  to 
produce  a  precipitate,  and  it  must  then  be  treated  ac- 
cording to  (79). 

Test  the  filtrate  for  metals  of  Groups  IV.,  III.,  II., 
and  I.  (See  page  115,  X.,  etc.) 


VII.— SOLUBILITY  OF  THE    SULPHYDRIC  ACID 
PRECIPITATE  IN  AMMONIC  SULPHIDE. 

Add  ammonic  sulphide  to  a  small  quantity  of  the 
precipitated  sulphides  (76),  and  warm  gently. 

If  the  precipitate  dissolves  entirely,  it  consists  of  the 
sulphides  of  metals  of  Group  VI.  Those  of  Group  V. 


METALS  OF  GROUP  VI. 


109 


are  absent.     Test  the  remainder  of  the  precipitate  accord- 
ing to  (82)  (a). 

(80)  If  a  part  of  the  precipitate  does  not  dissolve,  add 
four  or  five  parts  of  water,  and  separate  the  solution 
by  nitration  from  the  undissolved  precipitate. 

The  part  of  the  precipitate  which  is  insoluble  in  ammo- 
nic  sulphide,  after  being  carefully  washed,  must  be  tested 
for  sulphides  of  metals  of  Group  V.  (See  page  112, 
IX.) 

(81)  The  ammonic  sulphide  solution  obtained  in  (80) 
may  contain  metals  of  Group  VI.    Add  to  it  gradually 
dilute  chlorhydric  acid  until  the  solution  becomes  acid, 
and  observe  the  color  and  general  appearance  of  the 
precipitate  which  is  produced.     It  is  well  to  boil  the 
liquid  after  the  formation  of  a  precipitate. 

If  only  a  fine  white  precipitate  forms,  which  remains  a 
long  time  in  suspension  in  the  liquid,  even  after  boiling,  it 
consists  of  sulphur,  and  metats  of  Group  VI.  are  absent, 
and  the  tests  described  in  VIII.  can  be  omitted. 

A  flocculent  precipitate,  or  one  that  becomes  so  on  boil- 
ing, indicates  the  presence  of  metals  of  Group  VI.,  and 
the  color  of  the  precipitate  shows  what  metals  predomi- 
nate. Pass  to  the  following  tests : 

VIII— SEPARATION  OF  METALS  OF  GROUP  VI. 

(82)  If  the  test  (page  108,  VII.)  has  shown  the  presence 
of  metals  of  Group  VI.,  and  if  the  precipitate  with 
sulphydric  acid  (77)  was  not  entirely  soluble  in  am- 
monic sulphide,  the  whole  of  that  precipitate  must  be 
treated  two  or  three  times  with  ammonic  sulphide,  as 
directed  (79),  (80),  and  (81) ;  and  the  sulphides  of 
Group  VI.  must  be  precipitated  from  the  solution,  and, 
after  careful  washing,  treated  as  described  in  (83). 

(a)        Were  the  sulphides,  precipitated  by  sulphydric  acid, 


HO  PART  III. 

wholly  soluble  in  ammonic  sulphide  (see  79),  it  is 
sufficient  to  wash  and  dry  the  portion  of  the  precipi- 
tate (77)  which  was  not  treated  with  ammonic  sul- 
phide, and  to  use  it  for  (83)* 

(83)  Free  the  precipitate  (82)  as  completely  as  possible 
from  water  by  pressing  the  filter  and  its  contents  be- 
tween several  thicknesses  of  filter-paper,  remove  the 
precipitate  from  the  filter  and  heat  it  with  concen- 
trated chlorhydric  acid.  The  sulphides  of  antimony 
and  tin  will  be  dissolved,  while  the  sulphide  of  arsenic 
remains  undissolved.  Collect  the  residue  in  a  filter, 
wash,  and  dry  it  at  100°,  and  test  for  arsenic  as  fol- 
lows : 

Arsenic.  Place  a  small  portion  of  the  dried  residue  in  a 
(84:)  small  tube  closed  at  one  end,  and  put  over  it  about  six 
times  its  bulk  of  a  mixture  of  equal  parts  of  sodic  car- 
bonate and  potassic  cyanide.  Heat  the  portion  of  the 
tube  above  this  mixture,  and  afterwards  the  mixture 
itself,  gently  ;  if  as  the  result  of  this  any  moisture  is 
deposited  in  the  upper  part  of  the  tube,  wipe  it  out 
carefully  with  a  rolled-up  strip  of  filter-paper.  When 
the  whole  is  thoroughly  dry,  heat  the  lower  part  of  the 
tube  with  its  contents  to  a  red  heat.  ARSENIC,  if 
present,  is  sublimed  and  deposited  on  a  black  or 
brownish  ring  in  the  upper  and  cooler  part  of  the 
tube. 

Antimony.    Concentrate  the  solution  of  the  other  two  sul- 
(85)  phides  that  was  filtered  from  the  arsenious  sulphide 

*  By  heating  the  dry  precipitate  in  a  glass  tube,  or  with  less  accuracy, 
by  heating  it  before  it  is  dry,  on  a  bit  of  glass  or  porcelain,  an  approximate 
test  may  be  made  (see  13)  ;  and  in  case  arsenic  sulphide  alone  is  indicated 
by  the  complete  volatility  of  the  precipitate,  this  test  is  conclusive,  and  the 
remaining  tests  in  the  separation  of  metals  of  Group  V.  may  be  omitted. 
The  test  has  little  value  except  when  the  pure  yellow  color  of  the  precipi- 
tate gives  rise  to  the  suspicion  that  only  arsenic  sulphide  is  present. 


METALS  OF  GROUP  VI.  m 

(S3)  in  an  evaporator,  put  a  small  piece  of  zinc  in  the 
concentrated  solution,  and  bring  the  edge  of  a  piece 
of  platinum  foil  in  contact  with  the  zinc  for  a  minute 
or  two.  If  ANTIMONY  is  present  the  portion  of  the 
platinum  immersed  in  the  liquid  will  be  stained  black 
by  a  thin  deposit  of  that  metal. 
Tin.  Put  the  rest  of  the  concentrated  solution  of  the  sul- 

(86)  phides  in  an  evaporator  with  more  zinc,  collect  the 
precipitated  black  flakes  that  may  appear  after  a  time 
on  a  filter,  wash  carefully,  and  pour  a  little  concen- 
trated chlorhydric  acid  on  the  filter.     TIN,  if  present, 
will  be  dissolved,  and  the  solution  that  passes  through 
will  give  a  white  precipitate,  or  perhaps  a  gray  one  if 
much  metal  is  present,  with  mercuric  chloride.     (See 
page  56.) 

Gold.     (The  analytical  chemist  usually  knows  whether  it  is 

(87)  necessary  to  test  for  gold  or  not.    In  ordinary  analyses 
its  presence  would  be   improbable.)     When   gold  is 
present,  it  will  be  found  accompanying  the  sulphide  of 
arsenic  remaining  after  treatment  with  hot  chlorhydric 
acid  (83).     Heat  a  portion  of  this  residue  in  a  por- 
celain crucible  or  on  a  piece  of  a  broken  evaporator 
until  the  arsenic  and  excess  of  sulphur  have  been  vol- 
atilized.    Dissolve  the   portion   which   remains   in  a 
mixture  of  chlorhydric  and  nitric  acids.     Evaporate 
the  solution  nearly  to  dryness,  dilute  with  water,  and 
add  ferrous  sulphate   solution.     The  formation  of  a 
brown  or  purple  precipitate  of  METALLIC  GOLD,  either 
immediately  or  after  heating,  indicates  the  presence 
of  the  metal. 


IT2  PART  III. 

IX.— SEPARATION  OF  METALS  OF  GROUP  V. 
(SECTION  II.) 

(88)  If  a  portion  or  the  whole  of  the  precipitate  obtained 
with  sulphydric  acid  is  insoluble  in  ammonic  sulphide 
see  (80),  free  it  by  careful  washing  from  the  liquid  in 
which  it  was  formed,  or  from  the  ammonic  sulphide 
which  was  used  to  dissolve  the  soluble  portion,  place 
it  in  a  porcelain  dish,  pour  upon  it  pure  concentrated 
nitric  acid,  and  heat  it  gently,  if  red  fumes  are  given 
off,  until  they  cease.     In  any  case,  complete  the  oper- 
ation by  adding  a  little  water,  and  boiling  the  contents 
of  the  dish  for  a  few  minutes. 

If  no  part  of  the  precipitate ',  or  if  only  yellow  particles 
of  sulphur  remain  insoluble,  mercury  is  absent.     Pass  to 

(90). 

In  this  test,  when  the  liquid  holding  sulphur  in  sus- 
pension is  boiled,  the  sulphur  melts,  and  may  enclose 
particles  of  black  sulphide,  which,  then  become  very 
difficult  to  dissolve,  and  the  appearance  of  the  sulphur 
may,  in  such  a  case,  lead  to  an  erroneous  conclusion 
that  mercuric  sulphide  is  present.  It  is  for  this  reason 
that  the  precipitate  is  oxiflized  with  strong  nitric  acid, 
as  far  as  possible  at  'a- temperature  below  its  boiling 
point,  before  the  sulphides  are  finally  boiled  with  a 
somewhat  weaker  acid.  The  same  cause  makes  the 
confirmatory  test  for  mercury  with  stannous  chloride 
(89)  necessary,  when  there  appears  to  be  an  insolu- 
ble black  sulphide. 
Mercury.  If  a  black  sulphide,  HgS,  remains  insoluble 

(89)  after  the  above  treatment  (88),  a  MERCURIC  SALT  is 
probably  present.     Confirmatory  test :  Boil  the  black 
insoluble  sulphide  with  chlorhydric  acid  and  a  little 
potassic  chlorate  in  a  porcelain  dish,  and  evaporate 
until  the  greater  part  of  the  acid  is  volatilized  ;  dilute 


METALS  OF  GROUP  V.  (SECTION  II.)  113 

with  water  (it  is  not  necessary  to  filter),  and  add  stan- 
nous  chloride.  A  white  precipitate  of  MERCUROUS 
CHLORIDE,  Hg2Cl2,  is  formed,  if  mercury  is  present. 

If  mercury  is  present,  dilute  a  few  drops  of  the 
nitric  acid  solution  of  the  sulphides  (S3)  with  water, 
and  add  sulphydric  acid.  If  a  black  or  brown  preci- 
pitate is  formed  the  solution  must  be  tested  for  LEAD, 
BISMUTH,  COPPER,  and  CADMIUM.  Pass  to  (90).  If 
the  precipitate  is  yellow,  cadmium  alone  is  present 
(93).  If  no  precipitate  is  obtained  all  of  these  four 
metals  are  absent.  Pass  to  page  115,  X. 
LectcL.  Add  a  few  drops  of  strong  sulphuric  acid  to  a  small 

(90)  portion  of  the  nitric  acid  solution  of  the   sulphides 
(88)  and  evaporate  until  dense,  white  fumes  of  sul- 
phuric acid  appear,  and  dilute  with  a   considerable 
quantity  of  water.      If  a  white  precipitate  forms,   it 

Consists  Of  SULPHATE  OF  LEAD,  PbSO4. 

The  test  can  be  made  more  delicate  by  adding  an 
equal  bulk  of  alcohol  to  the  solution  after  it  has  been 
diluted  with  water.  If  lead  is  discovered,  treat  the 
whole  of  the  nitric  acid  solution  of  the  sulphides  in 
the  same  way  ;  filter  and  use  the  filtrate  for  (91). 
^Bismuth.  Add  ammonic  hydrate  to  alkaline  reaction,  to 

(91)  the  nitric  acid  solution  of  the  sulphides,  or  to  the  fil- 
trate from  the  lead  precipitate,  if  lead  was  present. 
If  bismuth  is  present  it  is  precipitated  as  the  HYDRATE 
OF  BISMUTH,  Bi(HO)8,  white.     If  bismuth  is  present, 
filter  and  use  the  filtrate  for  (92)  and  (93). 

Copper.     If  copper  is  present  it  is  dissolved  by  the  ammo- 

(92)  nic  hydrate,  and  imparts  a  blue  color  to  the  solution. 
Cadmium.     If  the  ammoniacal  solution  obtained  in  (91) 

(93)  was  colorless,  copper  is  absent,  but  cadmium  may  be 
present.     In  this  case  add  sulphydric  acid.      A  yellow 
precipitate,  CdS,  indicates  CADMIUM. 

If  copper  is  present,  neutralize  the  blue  ammoniacal 
8 


114  PART  III. 

solution  with  chlorhydric  acid  and  add  sulphydric 
acid.  This  will  precipitate  both  CdS  and  CuS  ;  the 
former  is  soluble  in  hot  dilute  sulphuric  acid.  After 
washing  the  mixed  sulphides  treat  the  mass  with  hot 
dilute  sulphuric  acid  and  filter.  Cool  the  filtrate  and 
pass  sulphydric  acid  again  through  the  liquid.  If 
cadmium  is  present  there  will  be  a  yellow  precipitate  of 
cadmic  sulphide,  CdS. 


METALS  OF  GROUPS  IV.  AND  III. 


METALS  OF  GROUPS  IV.  AND  III. 


X.— AMMONIC  SULPHIDE  TEST. 

(94i)  To  a  small  portion  of  the  nitrate  from,  the  precipi- 
tate produced  by  sulphydric  acid,  or  to  a  portion  of 
the  original  solution,  if  no  precipitate  is  produced  in 
it  by  sulphydric  acid,  add  sufficient  ammonic  hydrate 
(free  from  carbonate)  to  make  the  reaction  alkaline, 
and  then,  whether  a  precipitate  is  formed  or  not,  add 
ammonic  sulphide.  If  the  solution  contains  no  chlor- 
hydric  acid,  it  is  necessary  to  add  a  small  quantity 
before  neutralizing  with  ammonic  hydrate. 

If  a  black  precipitate  is  formed,  it  may  contain  the 

SULPHIDES  OF  NICKEL,  NiS  J  COBALT,  CoS  ;  IRON,  FeS  ; 

MANGANESE,  MnS  ;  and  ZINC,  ZnS  ;  and  the  HYDRATES 
OF  ALUMINIUM,  A12(HO)6,  and  CHROMIUM,  Cr2(HO)6. 
If  the  precipitate  is  white,  flesh-colored,  or  light  green, 
it  can  only  consist  of  the  SULPHIDES  OF  MANGANESE 
AND  ZINC,  and  the  HYDRATES  OF  ALUMINIUM  AND 
CHROMIUM.  In  this  latter  case  omit  (97),  (98),  and 
(101). 

If  no  precipitate  is  formed  no  members  of  Groups  III. 
and  IV.  are  present.  Pass  to  page  123,  XII. 
(,9t>)  If  a  precipitate  was  formed  in  the  above  test  (94), 
treat  a  considerable  quantity  of  the  solution  in  the 
same  way,  heat  the  liquid,  filter  as  quickly  as  possible, 
and  wash  immediately  with  boiling  water,  until  the 
nitrate  has  no  longer  an  alkaline  reaction. 

The  filtrate  must  be  tested  according  to  page  123,  XII. 


IT6  PART  III. 

SULPHIDES  INSOLUBLE  IN  DILUTE   CHLOR- 
HYDRIC  ACID. 

(96)  Add  to  the  precipitate  (95)  cold,  dilute  chlorhy- 
dric  acid  ;  if  a  black  residue  is  insoluble,  it  consists  of 
the    SULPHIDE  OF   NICKEL   or   COBALT.      Filter   and 
examine  the  residue  on  the  filter  according  to  (#7) 
and  (98). 

The  filtrate  must  be  tested  according  to  (99). 
If  the  precipitate  dissolves  entirely,  or  if  only  a  white 
residue  is  left,  no  NICKEL  or  COBALT  are  present.     Pass 
to  (99). 
Cobalt*     Dissolve  a  portion  of  the  precipitate,  which  proved 

(97)  to  be  insoluble  in  cold,  dilute  chlorhydric  acid,  in  the 
borax  bead,  and  expose  the  bead  to  the  action  of  the 
outer  blowpipe  flame.     If  the  bead  is  blue,  COBALT  is 
present.     If  the  bead  is  brown,  NICKEL  is  present  in 
large  quantity.    The  bead  is  blue  even  when  more  nickel 
than  cobalt  is  present.     To  test  for  traces  of  cobalt  in 
a  nickel  bead,  detach  the  hot  bead  from  the  wire,  heat 
it  two.or  three  minutes  on  charcoal  in  a  good  reduc- 
ing flame,  remove  it  from  the  charcoal,  and  melt  it  on 
the  platinum  wire  in  the  reducing  flame.      Even  if 
only  traces  of  cobalt   are  present,  the  bead  will  be 
colored  blue. 

Nickel.     When  nickel  is  present,  and  when  it  is  nearly  free 

(98)  from  cobalt,  it  can  be  discovered  by  the  brown  color 
which  it  imparts  to  the  borax  bead,  and  the  test  is 
conclusive. 

To  discover  small  quantities  of  nickel  in  the  pres- 
ence of  considerable  quantities  of  cobalt,  dissolve  the 
precipitate,  insoluble  in  cold  dilute  chlorhydric  acid 
(see  90),  in  concentrated  nitric  acid,  and  neutralize 
with  sodic  carbonate.  Potassic  cyanide  is  added  till 
the  resulting  precipitate  has  dissolved,  and  then  sodic 


METALS  OF  GROUPS  IV.  AND  HI.  II7 

hypochlorite  till  the  liquid  smells  strongly  of  it,  even 
after  being  shaken.  It  is  then  boiled.  If  nickel  is 
present,  a  black  precipitate  of  the  sesquioxide  (Ni2O3) 
is  obtained. 


SULPHIDES  SOLUBLE  IN  DILUTE  CHLORHYDRIC 

ACID. 

(00)  Boil  in  an  evaporating  dish  the  solution  which  was  ob- 
tained by  the  treatment  of  the  ammonic  sulphide  pre- 
cipitate with  cold,  dilute  chlorhydric  acid  (see  06), 
until  the  smell  of  sulphydric  acid  has  entirely  disap- 
peared ;  add  a  few  drops  of  strong  nitric  acid,  and  boil 
a  minute  longer,  and  filter  if  a  precipitate  of  sulphur 
has  formed. 

If  oxalic,  phosphoric,  and  boracic  acids  are  not  known 
to  be  absent,  pass  to  page  119,  XI. 

(100)  Add  sodic  hydrate  to  the  solution  obtained  in  (00) 
until  the  reaction  becomes  very  strongly  alkaline,  dilute 
with  water  and  boil  for  a  few  minutes. 

If  a  precipitate  forms  immediately,  or  after  boiling, 
test  it  according  to  (101)  and  (102)  ;  it  may  con- 
tain the  HYDRATES  of  IRON,  Fe2(HO)6,  MANGANESE, 

Mn(HO)2,  and  CHROMIUM,  Cr2(HO)6.  In  this  case 
filter  and  test  the  filtrate,  which  may  contain  ZINC  and 
ALUMINIUM,  according  to  (103)  and  (104). 

If  no  precipitate  forms,  pass  to  (103).     The  solu- 
tion may  contain  ZINC  and  ALUMINIUM. 
Iron.     Put  a  small  quantity  of  the  precipitate  (100)  on  a 

(101)  watch-glass,  dissolve  it  in  a  single  drop  of  dilute  chlor- 
hydric acid,  dilute  with  water,  and  add  potassic  sul- 
phocyanate.     A  red  color  indicates  the  presence  of 
IRON. 

Manganese  and  Chromium.    Fuse  a  portion  of  the 

(102)  precipitate  obtained  in  (100)  on  platinum  foil  with 


n8  PART  III. 

sodic  carbonate  and  sodic  or  potassic  nitrate.  A  green 
mass  indicates  MANGANESE,  owing  to  the  formation  of 
sodic  manganate  ;  a  yellow  mass,  CHROMIUM,  from  the 
formation  of  the  sodic  chromate.  A  very  small  quan- 
tity of  manganese  can  readily  be  detected  in  the  pres- 
ence of  a  considerable  quantity  of  chromium  by  the 
green  color  imparted  to  the  mass.  In  case  of  doubt, 
however,  place  the  platinum  foil  in  an  evaporator,  cov- 
er it  with  water,  add  a  few  drops  of  alcohol,  and  boil. 
The  sodic  chromate  will  be  entirely  dissolved,  along 
with  the  excess  of  sodic  carbonate,  while  the  sodic  man- 
ganate will  be  decomposed,  yielding  brown  flakes  of 
manganese  sesquioxide.  These  can  be  filtered  through 
a  small  filter,  washed  thoroughly  with  hot  water,  and 
fused  again  upon  platinum  foil,  as  before,  when  the 
green  color  will  appear  very  distinctly.  To  detect  a 
small  quantity  of  chromium  in  presence  of  excess  of 
manganese,  treat  the  mass  as  above  with  boiling  water 
and  filter.  The  appearance  of  a  yellow  color  in  the 
solution  is  proof  of  the  presence  of  CHROMIUM. 

Acetate  of  lead  may  also  be  added  to  the  solution, 
acidified  with  acetic  acid,  and  it  then  takes  a  deeper 
yellow  color  and  becomes  turbid,  or  a  yellow  precipi- 
tate is  formed  if  chromium  is  present.  The  formation 
of  a  white  precipitate  under  these  circumstances  only 
indicates  that  the  carbonate  of  sodium  employed  con- 
tained sulphate  of  sodium  as  an  impurity. 
Zinc.  Add  sulphydric  acid  to  a  portion  of  the  solution  in 
(103)  sodic  hydrate,  obtained  in  (100),  after  it  has  been 
filtered,  if  a  precipitate  was  formed.  If  a  white,  floccu- 
lent  precipitate  forms,  it  consists  of  SULPHIDE  OF  ZINC, 
ZnS. 

If  chromium  was  discovered  in  (102),  zinc  may 
also  be  present  in  the  precipitate  obtained  in  (100). 
Therefore,  in  that  case,  dissolve  a  portion  of  the  precip- 


METALS  OF  GROUPS  IV.  AND  III.  n9 

itate  by  boiling  with  a  very  little  dilute  chlorhydric  acid, 
add  sodic  hydrate  until  the  reaction  is  alkaline,  acidify 
with  acetic  acid  and  add  sulphydric  acid.  A  white, 
flocculent  precipitate  consists  of  SULPHIDE  OF  ZINC, 
ZnS. 

Aluminium.  Add  to  another  portion  of  the  sodic  hy- 
(104)  drate  solution,  obtained  in  (100),  chlorhydric  acid 
until  the  reaction  becomes  acid,  and  then  ammonic 
hydrate  until  it  becomes  alkaline,  and  boil.  If  a 
white,  flocculent  precipitate  forms,  it  consists  of  ALU- 
MINIC  HYDRATE,  A12(HO)6.  This  precipitate  is  at  first 
gelatinous,  and  it  may  easily  escape  notice  ;  it  is  there- 
fore best  to  set  the  test-tube  aside,  and  to  wait  a  quar- 
ter of  an  hour  for  the  precipitate  to  settle. 

XL— AMMONIC  SULPHIDE  TEST  IN  CASE  PHOS- 
PHORIC, OXALIC,  AND  BORACIC  ACIDS  ARE 
PRESENT. 

If  PHOSPHORIC,  OXALIC,  or  BORACIC  ACID  was  pres- 
ent  in  the  original  solution,  these  acids,  together  with 
the  metals  of  Group  II.,  may  be  contained,  wholly  or 
in  part,  in  the  solution  obtained  in  (99),  for  the  acids 
would  be  precipitated  with  any  of  the  metals  of  Groups 
II.,  III.,  and  IV.,  during  the  treatment  with  ammonic 
hydrate  and  ammonic  sulphide  in  (94),  and  the  pre- 
cipitates would  be  dissolved  during  the  treatment  with 
dilute  chlorhydric  acid  in  (96),  and  consequently  the 
metals  and  acids  might  be  contained  in  the  solution 
(99).  The  solution  (99)  must  be  tested  for  phos- 
phoric, oxalic,  and  boracic  acids,  and  freed  from  them 
before  the  ordinary  course  of  analysis  can  be  pro- 
ceeded with. 

PllOSpJioric  Acid.    Use  the  test  (127)  with  a  small 
portion  of  the  solution  obtained  in  (99). 


120  PART  III. 

If  phosphoric  acid  is  present,  use  (107)  and  the  suc- 
ceeding tests. 

If  phosphoric  acid  is  absent,  use  (100)  and  the  suc- 
ceeding tests. 

In  either  case  first  perform  the  operation  described  in 
the  next  paragraph. 

(103)  Before  testing  for  OXALIC  AND  BORACIC  ACIDS  it 
is  necessary  to  set  them  free  from  their  combinations 
with  the  metals  of  Groups  II.,  III.,  and  IV.  To  effect 
this  end  add  sodic  carbonate  to  a  small  quantity  of  the 
solution  obtained  in  (99)  until  the  reaction  becomes 
strongly  alkaline,  and  boil  for  a  few  minutes  and  filter. 
The  metals  are  precipitated,  with  the  exception  of  a 
portion  of  the  aluminium,  and  the  acids  remain  in  the 
solution. 

[If  no  precipitate  is  formed  with  sodic  carbonate  it 
*     is  unnecessary  to  test  further  for  these  metals  or  acids, 
as  in  that  case  they  cannot  be  present  in  the  solu- 
tion (99).] 

Oxalic  Acid.  Test  a  portion  of  the  filtrate  obtained  after 
boiling  with  sodic  carbonate  for  oxalic  acid  according 
to  (125). 

Itoracic  Acid.  Test  another  portion  of  the  same  filtrate 
for  boracic  acid  according  to  (128). 

If  one  or  both  these  acids  are  found \  the  whole  of  the 
remainder  of  the  solution  obtained  in  (99)  must  be  treated 
with  sodic  carbonate  as  in  (103),  and  the  precipitate  thus 
obtained  must  be  dissolved  in  dilute  chlorhydric  acid.  The 
solution  must  be  used  for  (107)  and  the  succeeding  tests, 
if  phosphoric  acid  was  discovered. 

If  phosphoric  acid  is  absent,  the  solution  must  be  used 
for  (100)  and  the  succeeding  tests. 

Aluminium.     If  it  was  found  necessary  to  treat  the  solu- 

(106)  tion  obtained  in  (99)  with  sodic  carbonate  according 
to  (105),  the  filtrate  from  the  precipitate  produced 


METALS   OF  GROUPS  IV.  AND  III.  I2I 

by  sodic  carbonate  may  contain  a  portion  of  the  alumin- 
ium. Add  to  the  nitrate  in  (103)  chlorhydric  acid 
until  the  reaction  becomes  acid,  and  then  ammonic 
hydrate  until  the  reaction  becomes  alkaline.  If  a 
white,  flocculent  precipitate  forms  immediately,  or  after 
long  standing,  it  contains  aluminium. 
Iron.  Add  a  few  drops  of  potassic  sulphocyanate  to  the 

(107)  solution  obtained  in  (99)  ;  a  red  color  indicates  the 
presence  of  IRON. 

(108)  Add  to  the  remainder  of  the  solution  obtained  in 
(99),  if  phosphoric'  acid  alone  is  present,  or  to  the 
solution  obtained  in  (105),  if  oxalic  or  boracic  acid 
is  likewise  present,  ferric  chloride,  until  a  few  drops, 
treated  with  ammonic  hydrate  on  a  watch-glass,  give 
a  yellow  and  not  &  white  precipitate  ;  dilute  largely  with 
water,   render   the   solution    alkaline   with   ammonic 
hydrate,  and  add  a  considerable  excess  ;   then  add 
acetic  acid  until  the  solution  has  a  slight  acid  reaction, 
and  boil  for  a  few  minutes  in  a  flask. 

The  precipitate  contains  all  the  IRON,  ALUMINIUM, 
CHROMIUM,  and  PHOSPHORIC  ACID  present  in  the  solu- 
tion so  treated.  Test  according  to  (109)  and 
(110). 

The  filtrate  contains  the  MANGANESE,  ZINC,  and 
probably  part  of  the  BARIUM,  CALCIUM,  and  MAGNE- 
SIUM which  were  present  in  the  original  solution. 

The  operations  described,  page  115,  X.,  #;z^page  123, 

XII.,  must  be  repeated  with  this  filtrate,  omitting  those 

which  relate  to  the  separation  of  NICKEL  and  COBALT, 

and  the  detection  of  IRON,  CHROMIUM,  and  ALUMINIUM. 

Chromium.     Test  a  portion  of  the  precipitate  obtained 

(109)  in  (108)  for  chromium  according  to  (102). 
Aluminium.     Boil  the  remainder  of  the  precipitate  ob- 

(110)  tained   in  (108)  with   sodic   hydrate,  and   test  the 
solution  for  aluminium  according  to  (104). 


I22  PART  III. 

[The  method  of  precipitation  by  boiling  the  acetic 
acid  solution  used  in  (10S)  can  be  used  in  all  cases 
for  the  separation  of  aluminium,  chromium,  and  iron 
(ferric  salts)  from  the  metals  of  all  other  groups,  except 
the  Groups  V.  and  VI.,  and  it  is  preferable  to  any  other, 
but  it  demands  more  skill  in  manipulation.] 


METALS  OF  GROUP  II. 


METALS  OF  GROUP  II. 

XII.— DETECTION  OF  BARIUM,  STRONTIUM,  CAL- 
CIUM, AND  MAGNESIUM. 

Should  the  filtrate,  after  removal  of  the  metals  of  Groups 
III.  and  IV.,  have  a  brown  color,  it  can  only  come  from  the 
presence  of  nickel,  a  small  quantity  of  the  sulphide  of  that 
metal  having  been  dissolved  in  the  excess  of  ammonic  hy- 
drate and  ammonic  sulphide  used.  Before  proceeding  to  the 
detection  of  metals  of  Group  II.,  the  nickel  must  be  entirely 
removed,  and  this  can  be  readily  accomplished  by  boiling  for 
a  few  minutes  and  filtering  again. 

(Ill)  To  a  small  portion  of  the  solution,  to  which  the 
previous  tests  have  been  applied,  or  to  a  solution  which 
has  been  found  to  contain  no  metals  of  the  higher 
groups,  add  ammonic  hydrate  until  the  reaction  be- 
comes alkaline,  and  then  ammonic  carbonate,  and  boil. 
(If  the  solution  does  not  already  contain  ammonic 
chloride,  this  also  must  be  added,  to  prevent  the  precip- 
itation of  magnesium.) 

If  a  white  precipitate  forms,  it  can  only  consist  of 

BARIC    CARBONATE,  BaCO3,     STRONTIUM    CARBONATE, 

SrCO8,  and  CALCIC  CARBONATE,  CaCO3. 

If  no  precipitate  forms,  BARIUM,  STRONTIUM,  and 
CALCIUM  are  absent ;  pass  to  (116). 

If  a  precipitate  was  formed  with  ammonic  carbonate, 
the  whole  of  the  solution  must  be  treated  as  described 


I24  PART  In' 

above.  Filter,  wash  the  precipitate,  and  test  the  fil- 
trate according  to  (116).  Dissolve  the  precipitate  by 
pouring  a  very  little  dilute  chlorhydric  acid  on  the 
filter,  and  use  the  solution  thus  obtained  for  (112), 
(113),  (114),  and  (115). 
Barium.  To  a  small  portion  of  the  solution  in  chlorhydric 

(112)  acid   add  a  considerable  quantity  of  calcic  sulphate. 
If  a  precipitate  forms  immediately,  it  consists  of  BARIC 
SULPHATE,  BaSO4. 

Strontium.     If   on  the   addition   of   calcic   sulphate,    as 

(113)  in  (112),  a  precipitate  appears  only  after  some  little 
time,  it  consists  of  STRONTIUM  SULPHATE,  SrSO4. 

( 1 14)  If  BARIUM  or  STRONTIUM  is  discovered  by  means  of 
calcic  sulphate,  add  dilute  sulphuric  acid  to  another 
portion  of  the  chlorhydric  acid   solution.     Boil,  filter, 
and  test  the  filtrate  for  calcium  (11&).     Examine 
the  precipitate  on  platinum  wire  moistened  with  chlor- 
hydric acid  in  the  flame.     (See  pages  26  and  34.) 
Barium  and  strontium  can  both  be  detected,  even  when 
a  small  quantity  of  one  is  present  with  a  large  quantity 
of  the  other.     After   placing  a  small  particle  of  the 
precipitate  on  the  loop  of  the  platinum  wire  the  parti- 
cle should  be    repeatedly   moistened   in   chlorhydric 
acid  and  subjected  again  to  the  action  of  the  flame. 
Where  strontium  is  present  in  very  small  proportion 
the  barium  color  will,  after  repeated  moistening  with 
chlorhydric   acid,  finally  give  place  to  the   crimson  of 
strontium.     Where,  however,  the  reverse  proportion  is 
found,  the  detection  of  barium  is  not  so  easy  by  this 
method.     In  this  case  add  to  a  portion  of  the  chlor- 
hydric acid   solution  obtained  in  (111)  a  solution  of 
strontium  sulphate.     A  faint  white  cloudiness  appear- 
ing after  some  time  indicates  the  presence  of  barium. 

Calcium.  If  barium,  or  strontium,  or  both,  were  present  in 

(115)  the  precipitate  obtained  on  the  addition  of  ammonic 


METALS  OF  GROUP  II.  125 

carbonate,  add  ammonic  hydrate  to  alkaline  reac- 
tion to  the  filtrate  obtained  after  precipitation  with 
dilute  sulphuric  acid  (114),  and  then  ammonic 
oxalate.  If  a  white  precipitate  forms,  it  consists  of 

CALCIC  OXALATE,  CaC2O4. 

If  neither  barium  nor  strontium  was  found,  the  whole 
of  the  precipitated  carbonate  (111)  must  have  con- 
sisted of  calcic  carbonate.  Confirm  by  flame  reaction, 
Page  35. 

It  is  quite  possible  that  owing  to  the  presence  of  a 
large  excess  of  ammonic  chloride  in  the  solution,  the 
calcium  may  have  escaped  precipitation  by  ammonic 
carbonate.  In  this  case  it  would  come  down  as  a  floc- 
culent  precipitate  in  the  test  for  magnesium  (116). 
To  detect  calcium  under  these  circumstances  filter 
the  precipitate  obtained  in  (110)  and  dissolve  it  in 
acetic  acid,  dilute  considerably  and  add  ammonic 
oxalate.  If  calcium  was  contained  in  the  precipitate 
-  obtained  by  hydric  disodic  phosphate,  it  will  now  be 
precipitated  as  the  oxalate,  which  is  insoluble  in  acetic 
acid.  Filter  the  calcic  oxalate,  and  to  the  filtrate  add 
ammonic  hydrate  to  alkaline  reaction,  when  the  ammo- 
nio-magnesic  phosphate  will  be  re-precipitated. 
Magnesium.  To  the  filtrate  from  the  precipitate  pro- 
(116)  duced  by  ammonic  carbonate  (111),  or  to  the  solu- 
tion in  which  no  precipitate  was  obtained  on  addition 
of  that  reagent,  add  sodic  phosphate.  If  a  white  pre- 
cipitate forms  (frequently  only  after  the  lapse  of  some 
minutes),  it  consists  of  the  AMMONIO-MAGNESIC  PHOS- 
PHATE, MgNH4PO4. 


126  PART  III. 


METALS  OF   GROUP  I. 

XIII.— DETECTION  OF  SODIUM,  POTASSIUM,  AND 
AMMONIUM. 

Before  testing  for  sodium  and  potassium,  precipitate  the 
metals  of  Groups  V.  and  VI.  which  are  present  with  sulphy- 
dric  acid,  and  precipitate  those  of  Groups  II.,  III.,  and  IV. 
with  a  mixture  of  ammonic  carbonate  and  sulphide,  and  use 
the  solution,  freed  from  those  metals,  for  the  following  tests 
(117)  and  (118): 
Sodium.  Evaporate  the  solution  to  dryness,  and  drive  off 

(117)  ammonia  salts,  if  any  are  present,  by  heat.     Sodium, 
if  present,  can  be  distinguished  by  the  yellow  color 
which  a  small  quantity  of  the  solid  residue  held  in 

•  the  flame  of  a  gas  or  alcohol  lamp  imparts  to  it.     See 
page  26. 

Frequently  sodium  can  be  detected  in  a  solution 
without  evaporation,  by  dipping  a  platinum  wire  in  the 
solution,  and  then  holding  it  in  the  flame. 
Potassium  can  be  recognized  by  the  violet  color  which  it 

(118)  imparts  to  the  flame.     The  solid  residue  obtained  in 
(117)  can  be  tested  for  potassium  in  the  same  way 
that  it  is  tested  for  sodium. 

If  sodium  is  also  present,  its  greater  coloring  power 
will  obscure  the  potassium  flame,  but  by  looking 
through  a  piece  of  blue  glass  at  the  flame,  the  violet 
color  can  be  distinguished  even  when  sodium  is  pres- 
ent. (The  color  of  the  potassium  flame  is  almost  the 
same  as  that  of  the  heated  wire,  while  the  sodium 


METALS  OF  GROUP  I.  127 

flame  is  much  more  blue,  if  it  is  not  excluded  entirely 
by  the  glass.) 

Add  to  a  portion  of  the  original  solution  a 
(HO)  few  drops  of  sodic  hydrate,  and  heat. 

AMMONIA  can  be  recognized  by  its  smell,  or  by  hold- 
ing a  piece  of  moistened  turmeric  paper  or  red  litmus 
paper  at  the  mouth  of  the  test-tube,  taking  care  not  to 
let  it  touch  the  sides,  which  may  be  moistened  with 
sodic  hydrate.  The  turmeric  paper  will  be  turned 
brown,  and  the  litmus  paper  blue,  if  ammonia  is 
present. 


I28  PART  III. 


TESTS  FOR  ACIDS. 


IT  is  usual  to  take  a  fresh  quantity  of  the  solution  to  test 
for  acids,  and  generally  the  tests  for  metals  precede  those  for 
acids,  in  order  that  information  gained  by  the  first  series  of  ex- 
periments may  point  out  the  most  convenient  way  of  detecting 
the  acids.  Silicic  acid,  however,  is  always  first  precipitated 
from  a  solution  as  directed,  page  101  (64),  and  the  acids  no- 
ticed under  the  chlorhydric  acid  test  (page  106,  V.)  are  to  be 
looked  for  while  that  test  is  applied  to  the  detection  of  the 
metals.  Phosphoric,  oxalic,  and  boracic  acids  interfere  with 
the  ordinary  methods  of  testing  for  the  metals  of  Groups  II., 
III.,  and  IV.,  and  consequently  these  acids  must  be  tested  for 
as  directed,  page  119,  XL,  during  the  application  of  the  tests 
for  the  metals  of  those  groups. 

In  all  other  cases,  when  the  presence  of  any  metals  or  acids 
in  the  solution  interferes  with  the  performance  of  the  test  for 
an  acid,  directions  are  given  under  the  head  of  each  acid  for 
removing  them. 

It  is  obvious  that  metals  and  acids  which  precipitate  each 
other  cannot  be  present  together  in  a  solution,  and  that  when 
certain  metals  have  been  found,  the  number  of  acids  which  it 
becomes  necessary  to  look  for  is  restricted  within  limits  de- 
termined by  this  consideration.  Therefore  the  knowledge 
already  acquired  of  the  composition  of  a  solution  must  be 
brought  to  bear  upon  the  problem  of  testing  for  acids. 

First,  the  reaction  which  the  solution  gives  with  test-paper 
must  be  considered,  and  then  the  tables  IV.  and  V.  must  be 


ACIDS  OF  GROUP  I. 


129 


consulted  to  ascertain  what  acids  can  exist  in  a  solution  pos- 
sessing the  observed  reaction,  together  with  the  metals  which 
have  been  discovered. 

For  instance,  if  a  solution  contains  lead,  and  is  neutral  or 
nearly  neutral,  the  only  acids  which  can  be  present  in  it  are 
acetic,  chlorhydric,  chloric,  and  nitric.  If  the  reaction  is 
strongly  acid  the  solution  may  contain  all  the  acids  except 
sulphuric,  sulphydric,  ferro-  and  ferricyanhydric.  Moreover, 
the  solution  cannot  contain  a  large  quantity  of  lead  and  chlor- 
hydric acid  at  the  same  time,  because  the  chloride  of  lead  is 
only  soluble  in  135  parts  of  water. 


ACIDS  OF  GROUP  L 

Arsenious  and  Arsenic  Acids  are  always  detected 

(120)  by  the  sulphydric  acid  test  in  searching  for  the  metals. 
When  these  acids  are  discovered,  they  must  always  be 
precipitated  by  sulphydric  acid  before  testing  further. 

Chromic  Acid  cannot  be  present  in  a  solution  to  which 

(121)  SULPHYDRIC    ACID   or  AMMONic  SULPHIDE  has  been 
added  (see  page  64).     If  chromic  acid  is  present  in  a 
solution,  it  must  be  contained  in  the  precipitate  ob- 
tained with  baric  chloride  (122),  and  can  be  detected 
by  heating  a  small  portion  of  the  precipitate  in  the 
borax  bead.     If  CHROMIC  ACID,  H2CrO4,  is  present, 
the  bead  will  be  colored  green.     Chromic  acid  can 
often  be  recognized  by  the  yellow  color  which  it  im- 
parts to  solutions  which  contain  it,  and  by  the  yellow 
precipitate,    PbCrO4,   which   is   obtained    by   adding 
plumbic  acetate  to  the  neutral  or  slightly  acid  solu- 
tion. 

9 


I3o  PART  III. 

XIV.— BARIC   CHLORIDE  TEST. 

Baric  nitrate  and  nitric  acid  should  be  used  instead  of  baric 
chloride  and  chlorhydric  acid,  when  lead,  silver,  or  mercurous 
salts  have  been  discovered  in  the  solution. 

(122)  Put  a  piece  of  litmus  paper  in  the  solution,  and  if 
the  reaction  is  acid,  add  ammonic  hydrate,  drop  by 
drop,  until  it  becomes  slightly  alkaline.     If  a  precipi- 
tate is  formed  in  consequence,  add  dilute  chlorhydric 
acid  only  in  sufficient  quantity  to  dissolve  it.     Add 
baric  chloride,  and  if  a  precipitate  forms,  it  indicates 
the  presence  of  ARSENIOUS,  ARSENIC,  CHROMIC,  SUL- 
PHURIC,  SULPHUROUS,    OXALIC,    FLUORHYDRIC,   PHOS- 
PHORIC, BORACIC,  CARBONIC,  OR  SILICIC  ACIDS.       (Sili- 

cic  acid  cannot  be  present  after  the  operation  (04) 
has  been  performed). 

If  these  acids  are  absent,  pass  to  the  argentic  nitrate 
test  (page  132,  XVIIL). 

(The  baric  chloride  test  is  of  little  value  except 
when  the  substance  is  soluble  in  water  with  a  neutral 
or  slightly  alkaline  reaction,  or  in  case  the  reaction  is 
acid,  when  the  metals  of  Groups  II.,  III.,  IV.,  and  V. 
are  absent.  The  following  special  tests  are  more  ac- 
curate :) 
Sulphuric  Acid.  Acidify  (if  the  solution  is  not  already 

(123)  acid)  with  dilute  chlorhydric  acid,  in  considerable  ex- 
cess, and  add  baric  chloride  as  long  as  a  precipitate 
continues  to  form.     (Use  nitric  acid  and  baric  nitrate 
if  chlorhydric  acid  produces  a  precipitate.)     If  sul- 
phuric acid  is  present,  it  is  precipitated  as  BARIC  SUL- 
PHATE, BaSO4,  fine  white  powder.     The  solution  must 
not  be  heated  when  it  is  intended  to  use  the  filtrate 
for  the  next  test. 

Sulphurous  Acid.     To  the  nitrate  from  the  precipitate 

(124)  produced  by  baric  chloride,  or  to  the  acid  solution  to 


CALCIC  SULPHATE  TEST.  131 

which  baric  chloride  has  been  added  without  produc- 
ing a  precipitate,  add  potassic  dichromate,  and  boil. 
If  a  precipitate  forms,  it  consists  of  BARIC  SULPHATE, 
BaSO4,  produced  by  the  oxidation  of  SULPHUROUS 
ACID,  H2SO3,  contained  in  the  solution.  Usually  the 
CHLORHYDRIC  ACID  TEST  (7<>)  is  more  convenient 
and  sufficiently  accurate. 

XV.— CALCIC  SULPHATE  TEST. 

If  the  solution  contains  metals  which  are  precipitated  by 
sulphydric  or  sulphuric  acids,  they  must  first  be  removed  by 
adding  a  slight  excess  of  those  precipitants  and  filtering. 

Add,  if  the  solution  is  alkaline,  add  acetic  acid 
until  the  reaction  becomes  acid  ;  if  it  is  acid,  add  sodic 
hydrate  until  the  reaction  becomes  alkaline,  and  then 
add  acetic  acid  until  it  becomes  acid,  and  test  as  below 
for  oxalic  acid.  If  a  precipitate  forms  and  does  not 
dissolve  in  the  acetic  acid,  add  to  the  original  solution 
a  considerable  excess  of  sodic  carbonate,  boil,  fil- 
ter, add  to  the  filtrate  acetic  acid  until  its  reaction 
becomes  acid,  and  test  as  follows  for  oxalic  acid : 
Add  calcic  sulphate  in  considerable  quantity.  If  a 
precipitate  forms,  it  consists  of  CALCIC  OXALATE, 
CaCsO*,  white  powder.  Fluorhydric  acid  is  the  only 
other  acid  which  precipitates  calcic  sulphate  under 
these  circumstances,  and  as  the  precipitate  is  almost 
transparent  and  gelatinous,  it  cannot  easily  be  mistaken 
for  that  produced  by  oxalic  acid. 

Acid  can  only  be  present  in  alkaline  solu- 
(126)  tions  in  glass  vessels.  If  there  is  reason  to  suspect 
the  presence  of  this  acid,  add  calcic  chloride  and  am- 
monic  hydrate  to  the  solution,  and  if  a  precipitate 
forms,  collect  it  on  a  filter,  and  examine  it  for  fluorine 
according  to  (44). 


132  PART  III. 

XVI.— AMMONIC  MOLYBDATE  TEST. 

Phosphoric  Acid.     Make  the  solution  strongly  acid  (if 

(127)  it  is  not  so  already)  with  nitric  acid,  and  add  a  small 
portion  of  it  to  a  considerable  quantity  of  ammonic 
molybdate  solution.*     If  phosphoric  acid  is  present, 
PHOSPHO-MOLYBDATE  OF  AMMONIUM,  yellow  crystalline 
powder,  is  precipitated.     If  the  quantity  of  phosphoric 
acid  in  the  solution  is  very  small,  the  precipitate  does 
not  form  until  after  several  hours. 

If  sulphydric  acid  is  present,  it  is  necessary  to  heat 
the  acid  solution  until  it  is  expelled,  before  perform- 
ing the  test.  See  (136)  (a)  for  this  test  in  the  pres- 
ence of  ferro-  or  ferricyanhydric  acid. 

XVII.— TURMERIC  PAPER  TEST. 

Boradc  Add.\     Strongly  acidify  the  solution  (if  it  is  not 

(128)  already  acid)  with  dilute  chlorhydric  acid,  dip  a  piece 
of  turmeric  paper  in  it,  and  dry  the  paper  by  holding 
it  over  the  lamp-flame,  without  charring  it.     If  a  red 
or  brownish-red  stain  appears  upon  the  paper  when  it  is 
dry,  it  is  due  to  the  presence  of  BORACIC  ACID,  H8BO8. 


ACIDS  OF  GROUP  II. 

XVIII.— ARGENTIC  NITRATE  TEST. 

Acidify  with  nitric  acid,  if  the  solution  is  not  already  acid, 
and  add  argentic  nitrate. 

Sulphydric  Acid.     If  a  black  precipitate  is  formed,  it 
(129)  must    contain    ARGENTIC    SULPHIDE,   Ag2S,    showing 

*See  foot-note,  page  67.  f  Ibid.,  page  68. 


ARGENTIC  NITRATE  TEST.  I33 

that  sulphydric  acid  was  present  in  the  solution.    The 
precipitate   may   also   contain   ARGENTIC   CHLORIDE, 

BROMIDE,      IODIDE,      CYANIDE,     FERRO-     and     FERRICY- 
ANIDE. 

If  a  precipitate  is  formed  which  is  not  black,  it  can 
only  be  due  to  presence  of  CHLORHYDRIC,  BROMHY- 

DRIC,    IODHYDRIC,     CYANHYDRIC,     FERROCYANHYDRIC, 
and  FERRICYANHYDRIC  ACIDS. 

If  no  precipitate  is  formed,  none  of  the  above  acids  are  present. 
Pass  to  (137). 

Chlorhydric  Acid.     See  also  (48).     Acidify  strongly 

(130)  with  concentrated  nitric  acid,  add  argentic  nitrate  in 
excess,  shake  thoroughly  if  a  precipitate  forms,  and 
allow  it  to  settle,  decant  the  liquid,  and  pour  on  the 
precipitate  strong  nitric  acid,  and  boil  for  five  minutes  ; 
if  a  precipitate   remains   undissolved,  it   consists  of 

ARGENTIC    CHLORIDE,     AgCl,    if    violet.       If  yellow,  of 

ARGENTIC  BROMIDE.     See  (132)  (a). 

See  (136)  (a)  for  this  test  in  the  presence  of  ferro- 
or  ferricyanhydric  acid. 
IBromhydric  Add  gives  with  argentic  nitrate  a  white 

(131)  curdy  precipitate,  which  darkens  on  exposure  to  light, 
is  insoluble  in  boiling  concentrated  nitric  acid,  and  is 
not  so  readily  soluble  in   ammonia   as  the   argentic 
chloride.     See  also  (49). 

lodliydric  Acid  gives  with  argentic  nitrate  a  yellow  pre- 

(132)  cipitate  of  the  iodide  which  is  nearly  insoluble  in  am- 
monia.    See  also  (50). 

(132)  Detection  ^/CHLORHYDRIC,  BROMHYDRIC,  and  IODHY- 
(d)  DRIC  acids  in  presence  of  each  other.  Mix  the  liquid  to 
be  tested  with  a  few  drops  of  dilute  sulphuric  acid, 
then  with  a  little  starch  paste,  and  add  a  few  drops  of 
fuming  nitric  acid,  or  a  solution  of  hyponitric  oxide 
in  sulphuric  acid,  when  the  blue  color  characteristic 


I34  PART  III. 

of  iodine  will  appear.  (See  page  74.)  Add  now 
chlorine  water  until  that  reaction  has  disappeared. 
On  continuing  the  gradual  addition  of  chlorine  water 
the  bromine  will  be  set  free,  and  will  impart  a  yellow 
or  brownish  color  to  the  liquid.  (See  page  73.) 
Chlorine  in  presence  of  bromine  can  best  be  detected 
by  the  following  method  :  Evaporate  the  solution  to 
dryness  and  mix  the  residue  with  a  little  potassic  di- 
chromate.  Place  the  mixture  at  the  bottom  of  a  clean 
test-tube  and  pour  on  it  a  few  drops  of  concentrated 
sulphuric  acid.  On  the  application  of  heat  dark  red 

drops    Of     CHROMYL     BICHLORIDE,    Or   CHROMIC    OXY- 

CHLORIDE,  CrO2Cl2,  will  condense  in  the  upper  por- 
tion of  the  test-tube.  Bromides,  under  the  same 
treatment,  give  a  similar  result ;  but  in  the  latter  case 
the  distillate  consists  of  bromine,  which  is  instantly 
.  decolorized  by  a  drop  of  ammonic  hydrate.  In  the 
case  of  chlorides,  when  the  CrO2Cl2  is  treated  with 
ammonic  hydrate,  it  gives  a  yellow  solution,  owing  to 
the  formation  of  ammonic  chromate  [(NH4)2CrOj. 

Where  the  substance  to  be  tested  for  bromine  and 
iodine  was  not  soluble  in  water  it  should  be  fused  on 
platinum  foil  with  sodic  carbonate.  The  mass  is  then 
treated  with  water  and  the  aqueous  solution  used  for 
the  foregoing  method  of  separation. 

XIX.— PRUSSIAN   BLUE   TEST   FOR   CYANHYDRIC 

ACID. 

Cyanhydric  Acid.  Add  to  the  solution  ferrous  sul- 
(133}  phate  and  a  few  drops  of  ferric  chloride  ;  add  sodic 
hydrate  until  a  precipitate  forms  (unless  the  solution 
is  alkaline  and  a  precipitate  forms  without  the  addition 
of  sodic  hydrate),  warm  for  a  minute,  and  add  dilute 
chlorhydric  acid  until  the  reaction  becomes  acid. 


FERRIC  CHLORIDE  AND  FERRO  US  SULPHA  TE  TESTS. 


135 


The  appearance  of  a  blue  precipitate  or  a  blue  color  in 
the  solution  is  evidence  of  the  presence'  of  cyanhydric 
acid. 

See  ( 136)  (b)  for  this  test  in  the  presence  of  ferro- 
or  ferricyanhydric  acid. 

XX.— FERRIC   CHLORIDE  TEST. 

FerrocyanTiydric  Acid.    Add  a  little  ferric  chloride 
(134=)  to  the  acid  solution.     If  ferrocyanhydric  acid  is  pres- 
ent, a  precipitate  of  PRUSSIAN  BLUE,  Fe4(FeCy6)3,  deep 
blue,  is  formed. 

XXI.— FERROUS   SULPHATE  TEST. 

Ferricyanhydric  Acid.    Add  a  little  ferrous  sulphate 
(135)  to  the  acid  solution.     If  ferricyanhydric  acid  is  pres- 
ent, a  precipitate  of  TURNBULL'S  BLUE,  Fe3(Fe2Cyi2), 
deep  blue,  is  formed. 

(130)  If  ferro-  or  ferricyanhydric  acid  is  present,  before 
(a)  performing  the  tests  for  PHOSPHORIC  ACID  (127),  and 
CHLORHYDRIC  ACID  (130),  the  following  steps  must 
be  taken :  Add  dilute  sulphuric  acid,  dilute  with 
water  if  the  solution  is  not  dilute,  add  cupric  sulphate, 
and  finally  add  enough  baric  nitrate  *  to  render  the 
precipitate  of  a  decidedly  lighter  color  ;  heat  almost 
to  boiling,  allow  the  precipitate  to  settle  for  a  few 
minutes,  filter,  and  use  the  filtrate  for  the  tests  (127) 
and  (130).  In  case  the  test  (51)  for  nitric  acid  is 
to  be  used  take  the  same  preliminary  steps  ;  using 
baric  chloride  in  place  of  baric  nitrate. 

*  Sulphuric  acid  and  baric  nitrate  or  chloride  are  only  added  in  order  to 
produce  a  heavy  precipitate  of  baric  sulphate,  which  carries  down  with  it 
the  lighter  particles  of  the  other  precipitates,  and  renders  the  filtration 
easier. 


I36  PART  III. 

(136)  If  ferro-  or  ferricyanhydric  acid  is  present,  the  test 
(b)  for  CYANHYDRIC  ACID  (133)  is  to  be  modified  in  the 
following  manner  :  Dilute  with  water  if  the  solution  is 
not  dilute,  add  dilute  sulphuric  acid,  then  add,  accord- 
ing as  ferro-  or  ferricyanhydric  acid  is  present,  ferric 
chloride  or  ferrous  sulphate,  or  both,  in  sufficient 
quantity  to  precipitate  the  ferro-  or  ferricyanhydric 
acid  or  both  acids  ;  finally,  add  baric  chloride  *  until 
the  color  of  the  precipitate  has  become  decidedly 
lighter  ;  shake  thoroughly,  allow  the  precipitate  to 
settle  for  a  few  minutes,  and  filter.  Add  to  a  portion 
of  the  filtrate  sodic  hydrate  until  a  precipitate  forms, 
warm  gently,  and  add  dilute  chlorhydric  acid  until 
the  solution  becomes  acid.  The  appearance  of  a  blue 
precipitate,  or  of  a  blue  color,  is  evidence  of  the  pres- 
ence Of  CYANHYDRIC  ACID. 


ACIDS  OF  GROUP  III. 

ACIDS  WHICH  ARE  NOT  PRECIPITATED  BY  ANY 
METALS. 

Chloric  Acid.    See  sulphuric  acid  test  (47). 

(13V) 
Nitric  Acid.  See  sulphuric  acid  test  (SI). 

(138) 
Acetic  Acid.  See  sulphuric  acid  test  (52). 

(139)  

*See  note,  page  135. 


INSOLUBLE  SUBSTANCES. 


CLASS  III. 

SUBSTANCES  WHICH  ARE  INSOLUBLE  IN  WATER 
AND  IN  ACIDS. 

(See  page  97.) 

The  only  substances  which  are  insoluble  after  the  treat- 
ment described  on  page  97  are  the  following  : 

Plumbic  Sulphate  (not  absolutely  insoluble  in  acids). 

Argentic  Chloride  (slightly  soluble  in  chlorhydric 
acid). 

Sulphur. 

Carbon. 

Baric  Sulphate,  Silica,  and  many  Silicates, 
and  some  Oxides. 

XXII.— SOLUTION   IN   AMMONIC    ACETATE    AND 
POTASSIC  CYANIDE. 

Plumbic  Sulphate.     Boil   a  portion  of  the  substance 
(14=0)  with  ammonic  acetate,  and  test  the  solution  (after  fil- 
tration,  if    necessary)    with   ammonic   sulphide.      If 
LEAD  is  present,  a  black  color,  or  a  black  precipitate  is 
formed. 

Test  the  solution  also  for  SULPHURIC  ACID,  accord- 
ing to  (123). 

If  lead  is  discovered,  repeat  the  treatment  with  am- 
monic acetate  until  no  more  lead  is  dissolved. 
Argentic  Chloride.    Digest  a  portion  of  the  substance, 
free  from  plumbic   sulphate  with   potassic    cyanide, 


138  PART  III. 

warm  (unless  it  blackens  by  warming),  and  test  the  so- 
lution (after  nitration,  if  necessary)  with  ammonic 
sulphide.  A  black  precipitate  indicates  the  presence 
of  SILVER.  If  a  black  precipitate  forms,  wash  it,  dis- 
solve it  in  strong  nitric  acid,  and  test  with  chlorhydric 
acid  according  to  (68)  in  order  to  confirm  the  pres- 
ence of  SILVER. 

If  silver  is  present,  repeat  the  treatment  with  potas- 
sic  cyanide  until  no  more  silver  is  dissolved. 
Sulphur.     Test  the  substance,  free  from  plumbic  sulphate 

(142)  and  argentic  chloride,  for  sulphur  according  to  (10). 
If  the  substance  is  moist,  it  must  be  carefully  dried  by 
heating  it  in  a  porcelain  dish  over  a  water-bath  before 
applying  the  test. 

If  sulphur  is  present,  heat  the  substance  in  a  cov- 
ered porcelain  crucible  until  the  sulphur  is  completely 
volatilized. 
Carbon.    If  the  substance  has  a  black  or  gray  color,  which  it 

(143)  loses  when  it  is  heated  with  the  blowpipe  on  platinum 
foil,  carbon  in   some  form  is  probably  present.     If 
carbon  is  present,  the  substance,  free  from  plumbic 
sulphate,  argentic  chloride,  and   sulphur,  should  be 
burnt,  until  as  much  as  possible  of  the  carbon  is  de- 
stroyed, by  heating  it  red-hot  on  platinum  foil  or  in  a 
porcelain  crucible. 


XXIII.— FUSION  WITH  POTASSIC  AND  SODIC 
CARBONATES  AND  SODIC  NITRATE. 

Baric  Sulphate,  Silicic  Acid,  and  many  Sili- 

(144)  cates,  and  some  Oxides.    Mix  the  finely  pow- 

(a)      dered  substance,  free  from  plumbic  sulphate,  argentic 

chloride,  and  sulphur,  and  as  nearly  free  from  carbon 

as  possible,  with  two  parts  of  potassic  carbonate,  two 


INSOLUBLE  SUBSTANCES.  139 

parts  of  sodic  carbonate,  and  one  part  of  sodic  nitrate  ;* 
bring  as  much  of  the  mixture  as  can  be  heated  at 
once  on  the  platinum  foil,  and  heat  the  under  side  of 
the  foil  with  a  blast-lamp  until  the  whole  mass  is  in  a 
state  of  quiet  fusion.  Repeat  this  operation  two  or 
three  times,  if  much  substance  is  required  for  the  ana- 
lysis. 

(b)  Detach  the  fused  mass  from  the  platinum  foil  each 
time  by  plunging  the  foil,  while  it  is  hot,  in  distilled 
water.  Boil  the  product  of  fusion  with  water,  and  if 
it  does  not  dissolve  completely,  filter,  and  wash  the 
precipitate  on  the  filter  with  distilled  water,  rejecting 
the  washings.  Continue  the  washing  until  baric 
chloride  ceases  to  produce  a  precipitate  in  the  water 
which  runs  through  the  filter. 

(143)  The  first  filtrate  may  contain  ARSENIC  ACID,  see 
(120),  (its  occurrence  is  rare)  ;  CHROMIC  ACID,  see 
(121)  ;  SULPHURIC  ACID  (123),  (the  tests  referred 
to  above  may  be  applied  successively  to  a  single  por- 
tion of  the  filtrate)  ;  FLUORHYDRIC  ACID  (its  occur- 
rence is  rare),  see  (126)  and  (44),  and  PHOSPHORIC 
ACID  (127).  The  two  last  tests  may  be  applied  suc- 
cessively to  another  portion  of  the  filtrate.  No  com- 
pound of  these  acids,  except  BARIC  SULPHATE,  is  by 
itself  insoluble,  but  insoluble  substances  sometimes 
contain  small  quantities  of  the  acids.  CALCIC  FLU- 
ORIDE is  only  decomposed  completely  by  the  treat- 
ment with  sulphuric  acid  described  in  (44). 
Silicic  Add*  The  principal  portion  of  the  filtrate  should 
(140)  be  tested  according  to  (04)  for  silicic  acid.  After 


*  The  sodic  nitrate  is  added  in  order  to  destroy  carbon  or  other  reduc- 
ing substances.  If  the  substance  to  be  analyzed  appears  to  contain  much 
carbon,  increase  the  quantity  of  sodic  nitrate.  If  the  substance  contains 
no  carbon,  the  use  of  sodic  nitrate  is  usually  unnecessary. 


140 


PART  III. 


separation  of  silica  the  only  metals  *  that  can  be  pres- 
ent in  the  chlorhydric  acid  solution  are  LEAD,  ALU- 
MINIUM, and  ZINC.  Test  for  lead  by  adding  an  excess 
of  dilute  sulphuric  acid  and  alcohol  to  the  solution. 
If  a  precipitate  of  PLUMBIC  SULPHATE  forms,  filter. 
Test  for  ALUMINIUM,  in  the  solution,  free  from  lead, 
by  adding  ammonic  hydrate  in  excess.  If  a  precipi- 
tate of  ALUMINIC  HYDRATE  forms,  filter.  Test  for 
ZINC  in  the  solution,  free  from  lead  and  aluminium  by 
adding  to  the  solution  containing  ammonic  hydrate  in 
excess  ammonic  sulphide.  A  flocculent,  white  precip- 
itate Consists  Of  SULPHIDE  OF  ZINC. 

(14f)  If  a  portion  remains  insoluble  after  boiling  the  fused 
mass  with  water  (1.44)  (&),  dissolve  it  in  chlorhydric 
acid.  If  much  silica  was  discovered  (see  146),  it 
is  best  to  evaporate  the  chlorhydric  acid  solution  to 
dryness,  and  to  proceed  as  directed  in  (64). 

Test  for  metals  in  the  chlorhydric  acid  solution  accord- 
ing to  page  107,  VI.,  etc. 

*  It  is  evident  that  sodium  and  potassium  in  insoluble  silicates  cannot 
be  detected  by  this  process.  All  reliable  methods  for  their  detection  re- 
quire the  use  of  platinum  vessels  and  great  care  in  manipulation.  Larger 
works  on  analysis  must  be  consulted  for  such  methods, 


EXPLANATION  OF   TABLES. 


THE  Tables  I.,  II.,  and  III.  contain  a  synopsis  of  the  course 
of  analysis  of  bodies  in  solution  given  in  Part  III.,  and  they 
are  intended  as  an  index  to  the  methods  which  are  there  de- 
scribed in  detail.  They  may  also  serve  as  guides  in  analytical 
work  to  students  who  have  made  themselves  acquainted  with 
the  detailed  descriptions  of  Part  III. 

A  skeleton  form,  similar  to  that  of  the  tables,  should  be  filled 
out  with  the  results  of  an  analysis,  and  the  reactions  which  oc- 
cur on  the  application  of  each  test  should  be  noted. 

The  sign  — ^ —  placed  under  the  formula  of  a  compound 
indicates  that  it  is  formed  as  a  precipitate  during  a  reaction. 
This  sign  is  used  in  the  following  tables,  and  it  will  also  be 
found  convenient  in  noting  the  results  of  analyses. 

The  Tables  IV.  and  V.  are  intended  to  indicate  the  degree 
of  solubility  in  water,  and  in  many  cases  in  alcohol,  acids,  and 
alkalies,  of  the  combinations  of  the  metals  and  acids  men- 
tioned in  Part  II. 

The  properties  of  a  salt  are  described  in  the  square  formed 
by  the  intersection  of  the  column  devoted  to  an  acid  with  that 
devoted  to  a  metal. 

The  Roman  numerals  standing  after  the  symbols  of  the 
metals  indicate  their  quantivalence,  and  the  formula  of  a  salt  is 
made  by  putting  the  symbol  of  a  metal  in  the  place  of  the 
symbol  of  an  equivalent  number  of  atoms  of  hydrogen  in  an 
acid.  When  an  acid  contains  more  than  one  atom  of  hydro- 
gen, several  classes  of  salts  may  be  formed,  according  as  one 
or  more  atoms  of  hydrogen  are  replaced  by  a  metal.  The 
normal  or  regular  salts  are  those  which  are  formed  by  the  re- 

141 


142 


PART  III. 


placement  of  the  greatest  passible  number  of  atoms  of  hydro- 
gen by  a  metal. 

The  descriptions  of  the  tables  refer  to  normal  salts,  but  the 
following  cases  are  exceptions,  because  the  salts  specified  are 
more  commonly  met  with  in  analysis  ;  and  in  using  the  tables 
the  formulas  below  must  be  substituted  for  those  of  the  normal 
salts : 

ARSENATES.— MgNH4AsO4  ;     MnNH4AsO4  ;    (Hg2)HAsO4 ; 

HgHAs04. 
PHOSPHATES.— (NH4)2HP04 ;    BaHPO4 ;   CaHPO4 ;  MgNH4 

P04 ;  MnNH4P04 ;  HgHPO4 ;  Na2HPO4. 
The  ARSENATE  OF  ALUMINIUM  probably  contains  more  acid 

than  the  normal  salt. 
The  CHROMATES  OF  ALUMINIUM  and  of  IRON  (ferric  chro- 

mate)  contain  a  larger  proportion  of  metal   than  the 

normal  salts. 
ARSENITES. — The  arsenites  referred  to  in  the  table  have  only 

two  atoms  of  hydrogen  replaced  by  a  metal,  except  Mg3 

(AsO3)2  and  Ag3AsO3. 
The  ARSENITES  OF  COBALT  and  MANGANESE  contain  less  than 

two  atoms  of  hydrogen  replaced  by  the  metal. 
BORATES—  (NH4)2B4O7 ;   BaB2O4 ;    CuB4O7 ;    FeB4O7 ;  (Fe2) 

B3O6 ;  PbB4O7 ;  CaB2O4 ;  MnB4O7 ;  NiB4O7 ;   K2B2O4  ; 
;  Na2B4OT;  ZnB4O7. 


EXPLANATION  OF  SIGNS  IN  TABLES  IV.  AND  V. 


Numbers   =  number  of  parts  of  water  required  to  dissolve  one  part  of  the 

anhydrous  salt  *  at  the  ordinary  temperature, 
oo  .  —  insolubility.!     The  sign  of  infinity  indicates  that  an  infinite 

quantity  of  water  is  required  to  dissolve  the  salt, 
s.  =  soluble  to  a  considerable  extent  in  water, 
s.s.  =  slightly  soluble, 
del.  =  deliquescent,  or  capable  of  dissolving  by  attracting  moisture 

from  the  air. 
dec.  =  decomposed.      Examples  :    dec.  =  decomposed    by    water. 

—  dec.  =  decomposed  by  acids. 
—  =  acids.     Example  :  —  s.  =  soluble  in  acids. 
+  =  sodic  or  potassic  hydrate.     Example  :  +  s.  =  soluble  in  so- 

dic  or  potassic  hydrate, 
am.  =  ammonic  hydrate, 
am.  cl.  =  ammonic  chloride. 

al.  =  alcohol.     Example  :  al.  00.=  insoluble  in  alcohol. 
When  no  solvent,  such  as  — ,  am.,  al.,  etc.,  is  indicated,  the  signs  co., 
s.,  s.s.,  and  dec.  refer  to  the  action  of  water  on  the  salt. 

*  The  salt  without  water  of  crystallization  is  referred  to. 

t  Most  of  the  salts  marked  insoluble  in  the  table  are  not  really  more  insoluble  than 
salts  like  baric  sulphate,  but  they  are  described  as  insoluble  because  they  are  known  to 
be  nearly  so,  and  because  the  quantity  of  water  required  to  dissolve  them  has  not  been 
determined. 

143 


fl 

.  * 

.fti 


SULPHYD 
Metals  of  G 


o 

' 


fi. 


1  1 


e 


yf 

fl 

2  I 


II 

l! 

•38 


K^ 
.So 


11! 


S^ 

§! 

.3 


ceo 


io 
io 
a 


28 
el 

f 

Is 


i£ 

b-o 

*8 

*  O 

••£  v 


ffl  « 


g  M  . 

'.2     „  C 
^O.2 

gWa 


O 


il 

u  o  3, 

111 


Ji 


) 


HNO 
dulate 
s  yello 


o  ^ 

3«, 

Sjl 
SI 

3^ 

W  ^ 
Ul 


ti 

§"l 


ilS. 


Me 


tl 


S< 


ll 


Silicic  Acid,  Si 

(64) 


id 

ens 


^E 
c-o 

il 


1° 

U-i    0 


Sulphurous  Ac 

(75)  gas  tu 
potassi 


II. 

bon- 
of 
ide. 
AG- 


OF  GROUP 
I. 


arbon 
als 
lor 
M 


m 
ini 
mm 
O3, 


f.tOil 

filffi. 


«1 


iil 
^"i 

S-r*c 

"S3  ^ 

1s3 
•a  .a  § 

S-o  4 

tf 

*SN 

Sl'fiJ 


l 


Ifw 

1M 

S*lH 

2l|  S 

lllS'l 
^^,0,-S 

^-•SsSl 


L 


o«- 

3§ 
$<$ 

WH 

SA 


«,_§:; 


color  of  fl 
metals  must 


lor  of 
a  blue  gl 
etals  mu 


col 
h 
me 


111 

^-5£ 

'8~ 


lution,  a 
H3,  ammonia 


ffc  .253 

b^ 


I 


OT 

a  1 
&  § 


p.  11 

ored 


^>  ^ 

w  1^ 

C  3^3 

-,  5" 


It 


,  freed  from  H^S  by  I 
,  is  boiled  with  NaH* 


§§ 

^    * 


'•J 

fi 


l 


s-g^ 


^8 


Mn(HO)2  imparts 
fused  on  platinum 


08'^° 


sg 


°0 


•so 


ll 

111 

o,-i  G 

a 

**•£ 
2+* 


Q  5 

3° 
P,S  = 


P 


C  NITRATE  TEST  :  p.  13 
of  Group  II.,  in  acid  solut 


•*•  1 

s 
§ 

a 

c 


5  2 

3! 

o     o 


93 


•£  «      * 

.•£.«         3 

rSM 


-'So     rt"-1 


* 


2S. 


I! 

•^ 
id 

9s 

& 


.2 

S-H 


y 
B 

o^ 


3« 

•d 


ol  11 


= 


4 

ui 

if 

8  I 

8  I 

1 

.  A 
8   1 

ti 

, 

.   ui 
8    1 

•'f 

- 

at 

d 

uf 
"13 

^T 

•t 

d 

3T 

T 

.   ui 

8  1 

in 

y 

C/J 

i" 

in 

8  "j' 

8?. 

8  I 

1! 
13 

8  I 

.  "" 

„• 

8I 

•j 

it 

"3 
•o 

d 

8- 
51 

8't 

in 

•?  J 

8*-o 
i 

ll 

1 

8 
I 

"f 

.£ 

! 

ui 

34 

d 

« 

-rt 

8 

*\ 

8f 

d 

d 

d 

u! 

d 

•  ^ 

. 

• 

P 

C 

• 

• 

a" 

CO^M 

8* 

8   in 

*? 

"w 

13 

.  in 

8   I 

«^ 

"3 

"O 

df 

8  T 

8  T 

8  T 

it 

d 

•O 

.  in 
8    1 

8*  f 

^ 

^ 

•  00 

8  1 

8  !! 

c35 

rt 

in 

| 

8 

H 

8 

1 

ui 

I 

1 

1 

1 

1 

in 

in 

8 

d 

M 

d 

5 

i 

gi 

8* 

in 

d"5 

in    | 

I'd 

.   in 
8     1 

in 

"rt 

"3 
•n 

•  in 
8    1 

8  6 

rt 

.  8 
8I 

8  P. 

8 

SI 

j  in 

^"rt 

T3"rt 

g 

in  .. 

8* 

ui 

i 

o 

"C  •£} 

"   1 

u 
8-S 
1 

s! 

ui 

I 

1 

1 

T 

! 

1 

• 

d 

3> 

. 

. 

. 

in 

,4 

1    * 

ui 

8  1 

8    1 

8 

8     1 

•S-a 

•S  1 

8    1 

d 

d 

"  13 

u? 

8! 

d 

8 

8  •; 

JM   fJJ 

| 

U 

3  « 

'0 

3d 

13 

•  f 

II 

j't 

13 
-o 

It 

^ 

•o 

ui 

in 

*i 

i 

If 

S 

00 

13 

it 

•  in 

8! 

i! 

J 

1 

it 

J 

ui 

d 

d 

ui 

B 

3  S 

1 

0 

| 

.  „; 

n." 

"! 

0 

1 

d 

•t 

d 

d 

"u-: 

d 

in 

in    j 

8  V." 

8  rf 

V) 

"rt 

v  ^_. 

o"- 
ft1" 

8  2 

u! 

^ 

8 

8', 

rt  CJ 

g 

F 

E 

"  F 

rt 

_ 

^ 

rt 

rt 

[ 

-C  -H 

a 

u; 

.  in 
8     1 

8  I 

i 

.  A 
8 

II 

It 

.  in 
8     | 

8  " 

.  8 
8 

•  8* 
8  1 

d 

1 

•o  I 

8  " 

8 

•o  1 

•1 

88 

.  u1 

8   1 

CCCQ 

EC 

S3' 

^8 

H    ^ 

. 

. 

2  in 

. 

d 

in   ™ 

d 

.in- 

«3 

1 

8- 

8  8 

,^ 

. 

3M 

"rt 

•*     1 

-    1 

A  | 

I 

" 

M 

•1 

"rt 

j,  | 

•n 

^•a 

a78 

1 

?' 

& 

"  1 

jj> 

4- 

0 

4- 

-j 

in 

80 

•Sd 

•o 

8 

•S 

•S 

8  1 

8 

5 

•o 

I 

8  T 
8 

ui   • 

1 

9  • 

>    t- 

si—  i 

•  in 

H 

3  M 

in 

IT 

. 

•i 

.  in 

r  « 

•  8 

8 

in 

ui 

ui 

„ 

o    •* 
SE 

zz, 

<c 

d 

in  u't 
"rt 

3  ~ 

fn 
1 

"rt 

"rt 

"rt 

in 

ui  U,J 

D^_. 

** 

"rt 

""•a" 

^ 

"rt 

S 

0- 

C- 

.2^ 

in 

in 

•  ui 

. 

in 

• 

•  in 

.   ui 

' 

.  y 

. 

'§-? 

,3 

8  1 

8  1 

*\ 

-O"rt 

8  1 

J-d 

•S-3 

8| 

8 

•S-a 

3 

8I 

^13 

t 

V 

8  ' 

Acetic  Acid,  HCaH3O2 

Arsenic  Acid,  H8AsO4 

Arsenious  Acid,  ) 
HaAs09.  (  ""I 

g 
!H 

1 

Bromhydric  Acid,HB.r 

Carbonic  Acid,  HaCOs 

ChlorhydricAcid,HCl. 

Chloric  Acid,  HC103. 

Chromic  Acid,HaCrO4 

CyanhydricAcid.HCy 

P 

^7 

tl 

!=: 

Fluorhydric  Acid,  HF. 

!2 

O 

•c 

1 
I 

Nitric  Acid,  HNO8... 

Oxalic  Acid,  H,CaO4  . 

Phosphoric  Acid,  ) 
H9P04.  (  -• 

Sulphuric  Acid,  HaSO4 

Sulphurous  Acid,) 
H,SOa  (•••" 

Sulphydric  Acid,  HaS. 

O 
ffi 

2 
1 

ieliquescent  ; 

^ 

c 

Nc 

N 

ui 

13 

8'I 

«1 

.  ui 
8    1 

ui 

r 

E 

N 

8f 

*S 

1 

ui 

I 

1 

ui 

I 

J 

ui  -: 

T 

3-3 

4 

E 

8'J 

1 

ui 
ui 

1 

.  8 

13 

Sl 

'o 

1 

II 

*>i 
•3 

^ 

ui 

•  ui 
8    1 

8  1 

ui 

.  £ 

'""  *rt 

8  1 

ui 

J 

y 

ui 

E 

u? 

'  ui 

8I 

ui 

ui 

88: 

I 

•§  •* 

H 

tn 

si 

ui 

<t 

.  ui 

8  1 

, 

.  ui 

8  1 

u? 

13 

it 

8 

8 

ui 

si 

u  ui 

•0    1 

ii 

•'l 

ui 

ui 

1 

¥,! 

n  « 

2  ^ 
si 

Strontium. 
Sr.  II. 

ei 

.  ui 

ui 

rf3 

13 

4 

ui 

*. 

am.cl.s^.;—  s 

~v 

•o 

_.  ui 

•o13 

T 

•a 

13 

ui 

ui 

10 

! 

T 

8^- 
y 

E 

rt 

1 

•j 

u,' 

^S 

5  i 
ii  1\ 

3     . 

^ 

^ 

ui 

x 

2  e| 

<• 

ui 

•a 

** 

N 

.  8 

13 

13 

A 

S 

8f 

ui 

•"•"' 

H  13 

^13 

N  j 

13 

r 

13 

ON 

ui    . 

13 

-  i 
*  n 

\* 

E 

ui 

J 

ui 

ui 

ui 

d 

E 

E 

c 

ui 

E 

tj 

^ 
II 

S3 

8 

ui 

1 

8  * 

ui 

1 

ri  rt 
I 

8i 

8  ^ 

ui 
1 

8  •: 
8 
1 

H 

8  1 

8  T 
8 

8  •; 

8 
1 

8  •; 
8 
1 

3 

13 

8 

•*•  <* 

T 

ui 

1 

8 

"T 

8  1 

11  i: 

il 

Potass'm 
K.  I. 

13  ^ 

Il3 

13 

ui 

10 

ii 

co"; 
"rt 

H13 

-' 

^1 

5J 

r: 

13 

^1 

ui 

*l 

ui 
co- 

"rt 

ui  8 

°Ha 

^1 

__.    ui 

of  solubility 
•  ammonic  Aj 

X* 

4 

ui 

I 

ui 
ui 

1 

.  ui 
8    1 

ui 

8t 

4 

T3 

ui 

.i 

s  •; 

8 

1 

ui 

E 

.  rt 
8  T 

8 
1 

i\ 

8 
I 

i\ 

15 

'U 

.  " 
13 

ui 
1 

8  1 

* 

8-S 
1 

8  1 

II 

I'l 

•Jr. 

1= 

10 
N 

.  ui 

8  1 

.  ui 

8  1 

& 
n 

8  I 

1O 

* 

.  ui 

8      1 

••! 

8 

W     J 

8 

sif 

•o  1 

.  ui 

8  I 

8  I 

IT' 

•3  | 

'•  8 

8I 

imbers  = 
=  alkalie 

&** 
SK 

* 

.  ui 

8  1 

.  ui 

8   1 

8 

.  ui 

8  1 

8 

.  ui 

8  g 

8 

•  8 
8   1 

8 

8 

•d  1 

8  I 

8  f 

i 

.8 
8  1 

•  + 

?^ 

tn 

P 

ui 

.  ui 

8  1 

.  ui 

8  1 

.  ui 
8    1 

•s 

•O 

.  ui 

8  1 

««* 

ui 

8I 

8I 

8 

8 

ui  rt^ 

8 

1 

.  ui 
8    1 

i 

13 

it 

8t 

10 

"T 

8T 

1 
I-jj 

S  1 
a  1 

I'B 

C  ti 

H 

ui 

ft 

.  ui 

8| 

if 

13 

8- 
S1! 

°°.  *? 

1 

ui 

, 

ui 

ui 

••? 

13 

ui 

13 

u> 

.  15 

8  E 

§  rt 

U? 
1 

if 

.  8 

ro    . 

13 

i-j 

y 

w'T 

wl 
Si! 

M 

|a 

4 

13 

.  ui 

8  1 

.  ui 

8  I 

•  ui 

8  1 

< 

o    • 

ll 

fff 

"5  "s 

rt 

5  uT 

81 

.  8 
8  1 

if 

ui-5 
ui    | 

i 

IO 

it 

.  ui 
8    1 

8 

1 

81 

•f 

TABLE  V.—  EXPLAN 

Acetic  Acid,  HCaH3O2 

Arsenic  Acid,  H3AsO4 

Arsenious  Acid,  ) 
H,As08.  S~~ 

Boracic  Acid,  H3BO3  . 

Bromhydric  Acid.HBr 

V^arbonic  Acid,  HaCO3 

^XfhlorhydricAcid.HCl. 

Chloric  Acid,  HC1OS. 

Chromic  Acid,H2CrO4 

Cyanhydric  Acid,  HCy 

•-U 

|| 

^ 

Fluorhydric  Acid,  HF. 

X 

.g 

O 

K 

y 

'B 

Oxalic  Acid,  H2C2O4  . 

y 

If 

N/Sulphuric  Acid,  H2SO4 

!2 

P" 

,/Sulphydric  Acid,  H2S. 

~1VO  'OOSIONVMd  NVS 


"  '!/!  .SH3AVSSV 

^!AiI^S 
(2H:  ilJLSflf 


.Vki     Ci'lOW 


14  DAY  USE 

RETURN  TO  DESK  FROM  WHICH  BORROWED 

LOAN  DEPT. 

This  book  is  due  on  the  last  date  stamped  below,  or 

on  the  date  to  which  renewed. 
Renewed  books  are  subject  to  immediate  recall. 


8  May'GSPSf 
REC'D  LP 

MAY  8   1963 

200ef65CD 


REC'D  LD 


OC1     6'65-9pM 


4  - 1968 


LD 


8EC1JLQ.  FEB  23  71 -2PM  5  7 


LD  21A-50m-8,'61 
(Cl795slO)476B 


General  Library 

University  of  California 

Berkeley 


M81896 


c.7 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


